LIBRARY 

THE  UNIVERSITY 
OF  CALIFORNIA 
SANTA  BARBARA 

PRESENTED  BY 

PROFESSOR  GEORGE  HAND 


ANTOINE      LAURENT     LAVOISIER 

FERMIER  GENERAL    NE  A  PARIS  LE  iff  AOUT 
16   ®>-<caX,    COM.  2  . 


ANTOINE  LAURENT  LAVOISIER  (1743-1794) 

Famous  for  his  care  in  quantitative  experiments,  for  demonstrating  the  true 
nature  of  combustion,  for  introducing  system  into  the  naming  and  grouping 
of  chemical  substances.  Executed  (1794)  during  the  French  Revolution  be- 
cause of  his  connection  with  the  government.  This  picture  is  taken  from  a 
French  engraving  of  1799.  The  panel  represents  Lavoisier  as  he  i 
arrested  in  his  laboratory  by  the  Revolutionary  Committee 


AN  ELEMENTARY  STUDY 
OF  CHEMISTRY 


BY 
WILLIAM  McPHERSON 

AND 

WILLIAM  EDWARDS  HENDERSON 

PROFESSORS    OF    CHEMISTRY,  OHIO    STATE    UNIVERSITY 


SECOND  REVISED  EDITION 


GINN  AND  COMPANY 

BOSTON     •    NEW  YORK     •     CHICAGO     •     LONDON 
ATLANTA     •     DALLAS     •     COLUMBUS     •     SAN    FRANCISCO 


COPYRIGHT,  1905,  1906,  1917,  BY 
WILLIAM  McPHERSON  AND  WILLIAM  E.  HENDERSON 


ALL   RIGHTS   RESERVED 
A  621.3 


gfte  gtftengnm 

GINN  AND  COMPANY  •  PRO- 
PRIETORS •  BOSTON  •  U.S.A. 


PREFACE  TO  THE  SECOND  EDITION 

The  advance  of  chemistry  in  the  last  decade  has  made 
necessary  the  revision  of  this  text.  While  thoroughly 
revising  the  subject  matter  the  authors  have  made  some 
alterations  in  the  order  of  presentation.  Carbon  and 
carbon  dioxide  are  presented  at  an  earlier  point ;  the 
chapter  on  neutralization  is  preceded  by  a  brief  chapter 
devoted  to  a  metal  (sodium)  and  a  base  (sodium  hydrox- 
ide), and  one  devoted  to  a  nonmetal  (chlorine)  and  an 
acid  (hydrochloric  acid)  ;  the  space  given  to  the  compounds 
of  carbon  has  been  extended  a  little,  and  the  material 
has  been  brought  forward  into  its  appropriate  place  in 
the  text. 

In  justification  of  this  last  change  the  authors  would 
call  attention  to  the  fact  that  a  very  large  percentage  of 
those  who  take  one  year  of  chemistry  do  not  continue  the 
subject.  It  seems  unreasonable  that  after  a  year  of  study 
the  student  should  have  no  knowledge  of  the  most 
important  of  the  organic  compounds,  for  he  will  meet 
with  them  in  everyday  life  far  oftener  than  with  the 
majority  of  the  compounds  of  inorganic  chemistry. 

The  other  changes  incorporated  in  this  revision  are  those 
which  are  suggested  by  the  development  of  the  science  in 
the  past  ten  years. 

The  authors  wish  to  express  their  indebtedness  to 
Mr.  E.  L.  Mahaffey  and  Mr.  J.  H.  Young  for  valuable 
assistance  rendered  in  the  revision  of  the  text. 

THE  AUTHORS 
OHIO  STATE  UNIVERSITY 
COLUMBUS,  OHIO 


PREFACE  TO  THE  REVISED  EDITION 

In  offering  this  book  to  teachers  of  elementary  chemistry 
the  authors  lay  no  claim  to  any  great  originality.  It  has 
been  their  aim  to  prepare  a  textbook  constructed  along 
lines  which  have  become  recognized  as  best  suited  to  an 
elementary  treatment  of  the  subject.  At  the  same  time 
they  have  made  a  consistent  effort  to  make  the  text  clear 
in  outline,  simple  in  style  and  language,  conservatively 
modern  in  point  of  view,  and  thoroughly  teachable. 

The  question  as  to  what  shall  be  included  in  an  ele- 
mentary text  on  chemistry  is  perhaps  the  most  perplexing 
one  which  an  author  must  answer.  While  an  enthusiastic 
chemist  with  a  broad  understanding  of  the  science  is  very 
apt  to  go  beyond  the  capacity  of  the  elementary  student, 
the  authors  of  this  text,  after  an  experience  of  many  years, 
cannot  help  believing  that  the  tendency  has  been  rather  in 
the  other  direction.  In  many  texts  no  mention  at  all  is 
made  of  fundamental  laws  of  chemical  action  because  their 
complete  presentation  is  quite  beyond  the  comprehension 
of  the  student,  whereas  in  many  cases  it  is  possible  to 
present  the  essential  features  of  these  laws  in  a  way  that 
will  be  of  real  assistance  in  the  understanding  of  the 
science.  For  example,  it  is  a  difficult  matter  to  deduce 
the  law  of  mass  action  in  any  very  simple  way;  yet  the 
elementary  student  can  readily  comprehend  that  reactions 
are  reversible,  and  that  the  point  of  equilibrium  depends 
vii 


viii     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

upon  rather  simple  conditions.  The  authors  believe  that 
it  is  worth  while  to  present  such  principles  in  even  an 
elementary  and  partial  manner  because  they  are  of  great 
assistance  to  the  general  student  and  because  they  make 
a  foundation  upon  which  the  student  who  continues  his 
studies  to  more  advanced  courses  can  securely  build. 

The  authors  have  no  apologies  to  make  for  the  extent 
to  which  they  have  made  use  of  the  theory  of  electrolytic 
dissociation.  It  is  inevitable  that  in  any  rapidly  developing 
science  there  will  be  differences  of  opinion  in  regard  to  the 
value  of  certain  theories.  There  can  be  no  question,  how- 
ever, that  the  outline  of  the  theory  of  dissociation  here 
presented  is  in  accord  with  the  views  of  the  very  great 
majority  of  the  chemists  of  the  present  time.  Moreover, 
its  introduction  to  the  extent  to  which  the  authors  have 
presented  it  simplifies  rather  than  increases  the  difficulties 
with  which  the  development  of  the  principles  of  the  science 
is  attended. 

The  oxygen  standard  for  atomic  weights  has  been 
adopted  throughout  the  text.  The  International  Com- 
mittee, to  which  is  assigned  the  duty  of  yearly  reporting 
a  revised  list  of  the  atomic  weights  of  the  elements,  has 
adopted  this  standard  for  their  report,  and  there  is  no 
longer  any  authority  for  the  older  hydrogen  standard. 
The  authors  do  not  believe  that  the  adoption  of  the  oxy- 
gen standard  introduces  any  real  difficulties  in  making 
perfectly  clear  the  methods  by  which  atomic  weights  are 
calculated. 

The  problems  appended  to  the  various  chapters  have 
been  chosen  with  a  view  not  only  of  fixing  the  principles 
developed  in  the  text  in  the  mind  of  the  student  but 
also  of  enabling  him  to  answer  such  questions  as  arise 


PREFACE  ix 

in  his  laboratory  work.  They  are,  therefore,  more  or  less 
practical  in  character.  It  is  not  necessary  that  all  of  them 
should  be  solved,  though  with  few  exceptions  the  lists  are 
not  long.  The  answers  to  the  questions  are  not  directly 
given  in  the  text,  as  a  rule,  but  can  be  inferred  from  the 
statements  made.  They  therefore  require  independent 
thought  on  the  part  of  the  student. 

With  very  few  exceptions  only  such  experiments  are 
included  in  the  text  as  cannot  be  easily  carried  out  by  the 
student.  It  is  expected  that  these  will  be  performed  by 
the  teacher  at  the  lecture  table.  Directions  for  laboratory 
work  by  the  student  are  published  in  a  separate  volume. 

While  the  authors  believe  that  the  most  important 
function  of  the  elementary  text  is  to  develop  the  princi- 
ples of  the  science,  they  recognize  the  importance  of 
some  discussion  of  the  practical  application  of  these  prin- 
ciples to  our  everyday  life.  Considerable  space  is  there- 
fore devoted  to  this  phase  of  chemistry.  The  teacher 
should  supplement  this  discussion  whenever  possible  by 
having  the  class  visit  different  factories  where  chemical 
processes  are  employed. 

Although  this  text  is  now  for  the  first  time  offered  to 
teachers  of  elementary  chemistry,  it  has  nevertheless  been 
used  by  a  number  of  teachers  during  the  past  three  years- 
The  present  edition  has  been  largely  rewritten  in  the  light 
of  the  criticisms  offered,  and  we  desire  to  express  our 
thanks  to  the  many  teachers  who  have  helped  us  in  this 
respect,  especially  to  Dr.  William  Lloyd  Evans  of  this 
laboratory,  a  teacher  of  wide  experience,  for  his  continued 
interest  and  helpfulness.  We  also  very  cordially  solicit 
correspondence  with  teachers  who  may  find  difficulties 
or  inaccuracies  hi  the  text. 


x         AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  authors  wish  to  make  acknowledgments  for  the 
photographs  and  engravings  of  eminent  chemists  from 
which  the  cuts  included  in  the  text  were  taken ;  to 
Messrs.  Elliott  and  Fry,  London,  England,  for  that  of 
Ramsay ;  to  The  Macmillan  Company  for  those  of  Davy 
and  Dalton,  taken  from  the  Century  Science  Series;  to 
the  L.  E.  Knott  Apparatus  Company,  Boston,  for  that 

of  Bunsen. 

THE  AUTHORS 
OHIO  STATE  UNIVERSITY 
COLUMBUS,  OHIO 


CONTENTS 

CHAPTER  PAGE 

I.   MATTER  AND  ENERGY 1 

II.    VARIETIES   OF   MATTER:    COMPOUNDS,   ELEMENTS, 

MIXTURES 12 

III.  OXYGEN.     ..".'.    .    .    .    .    .     .     .     .     ...  24 

IV.  HYDROGEN  .    .    .    .     .    .     :  '..    ...     .    .    .     .  38 

V.    THE  GAS  LAWS;    THE  KINETIC  THEORY    ....  52 

VI.   WATER  ;    HYDROGEN  PEROXIDE     ...'..;.  64 

VII.   COMBINING  WEIGHTS;   THE  ATOMIC  THEORY      .     .  86 

VIII.   FORMULAS  ;   EQUATIONS  ;   CALCULATIONS  ....  93 

IX.   THE  THREE  STATES  OF  MATTER 104 

X.   CARBON  AND  CARBON  DIOXIDE 116 

XI.    NITROGEN    AND    THE    RARE    ELEMENTS    IN    THE 

ATMOSPHERE 128 

XII.   THE  ATMOSPHERE    .     .....     .     .     .     .     .     .     .  135 

XIII.  SOLUTIONS;   THE  IONIZATION  THEORY 143 

XIV.  CHLORINE  ;  HYDROGEN  CHLORIDE  ;  HYDROCHLORIC 

ACID 159 

XV.   SODIUM;   SODIUM  HYDROXIDE  ........  172 

XVI.   ACIDS,  BASES,  AND  SALTS;    NEUTRALIZATION    .     .  179 

XVII.   VALENCE 193 

XVIII.   COMPOUNDS  OF  NITROGEN 200 

XIX.   SPEED  OF  REACTIONS;   EQUILIBRIUM 219 

XX.   SULFUR;   SELENIUM;   TELLURIUM 228 

XXI.   THE  PERIODIC  LAW 254 

XXII.   THE  CHLORINE  FAMILY 264 

XXIII.   MOLECULAR  WEIGHTS;   ATOMIC  WEIGHTS     .     .     .  279 


xii      AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

CHAPTER  PAGE 

XXIV.   CARBON  MONOXIDE  ;  CARBONIC  ACID  ;   HYDRO- 
CARBONS        291 

XXV.    FUELS  ;   FLAMES  ;   ELECTRIC  FURNACES    .     .     .  306 
XXVI.   CARBOHYDRATES;    ALCOHOLS;    COAL-TAR    COM- 
POUNDS      324 

XXVII.   ORGANIC  ACIDS;   FATS  AND  OILS 339 

XXVIII.   THE  PHOSPHORUS  FAMILY 346 

XXIX.   SILICON  ;   TITANIUM  ;   BORON 367 

XXX.   THE  COLLOIDAL  STATE 382 

XXXI.    THE  METALS 389 

XXXII.   THE  ALKALI  METALS 396 

XXXIII.  SOAP;    GLYCERIN;    EXPLOSIVES 416 

XXXIV.  THE  CALCIUM  FAMILY;  FERTILIZERS  ....  423 
XXXV.   THE  MAGNESIUM  FAMILY 440 

XXXVI.   THE  ALUMINIUM  GROUP 451 

XXXVII.   ALUMINIUM  SILICATES  AND  THEIR  COMMERCIAL 

APPLICATIONS 464 

XXXVIII.   THE  IRON  FAMILY 469 

XXXIX.   COPPER,  MERCURY,  AND  SILVER 493 

XL.    TIN  AND  LEAD 511 

XLI.   MANGANESE  AND  CHROMIUM 524 

XLII.   URANIUM  ;   RADIUM  AND  THORIUM 535 

XLIII.   THE  PLATINUM  METALS  AND  GOLD      ....  544 

XLIV.    SOME  APPLICATIONS  OF  THE  RARER  ELEMENTS  553 

APPENDIX 557 

INDEX  .                                                                                          .  563 


AN  ELEMENTARY  STUDY 
OF  CHEMISTRY 

CHAPTER  I 
MATTER  AND  ENERGY 

The  natural  sciences.  Before  we  advance  very  far  in 
the  study  of  nature  it  becomes  evident  that  the  one  large 
study  must  be  divided  into  a  number  of  more  limited  ones 
for  the  convenience  of  the  investigator  as  well  as  of 
the  student.  These  "more  limited  studies  are  called  the 
natural  sciences. 

Since  the  study  of  nature  is  divided  in  this  way  for 
mere  convenience  and  not  because  there  is  any  division  in 
nature  itself,  it  often  happens  that  the  different  sciences  are 
very  intimately  related  and  that  a  thorough  knowledge  of 
any  one  of  them  requires  a  considerable  acquaintance  with 
several  others.  Thus  the  botanist  must  know  something 
about  animals  as  well  as  about  plants ;  the  student  of 
human  physiology  must  know  something  about  physics 
as  well  as  about  the  parts  of  the  body. 

Intimate  relation  of  chemistry  and  physics.  Physics  and 
chemistry  are  two  sciences  related  in  this  close  way,  and  it 
is  not  easy  to  make  a  precise  distinction  between  them.  In 
a  general  way  it  may  be  said  that  they  are  both  concerned 
with  inanimate  matter  rather  than  with  living  matter,  and 
more  particularly  with  the  changes  which  such  matter 
may  be  made  to  undergo. 


2        AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Changes  in  nature.  The  thoughtful  observer  of  nature 
is  constantly  impressed  by  the  fact  that  everything  about 
him  is  undergoing  change.  All  living  organisms  pass 
through  a  cycle  of  birth,  development,  maturity,  decline, 
and  death.  Many  solid  rocks  disintegrate  rapidly  when 
exposed  to  the  action  of  the  weather;  with  others  the 
changes  are  too  slow  to  be  perceptible,  but  geologists 
have  been  able  to  show  that  even  the  most  permanent 
rocks  are  gradually  altered  by  long  exposure  to  air  and 
moisture.  The  metals  we  win  from  the  ores  in  time  rust 
and  corrode ;  the  structures  we  rear  crumble ;  the  electric 
current  we  generate  is  changed  into  the  motion  of  the  car, 
the  light  of  the  lamp,  and  the  heat  of  the  wires. 

Questions  suggested  by  the  burning  of  fuel.  Doubtless 
one  of  the  changes  that  earliest  attracted  the  attention 
of  man  was  that  which  is  evident  during  the  burning  of 
fuel.  If  we  set  fire  to  a  piece  of  coal  or  wood,  most  of 
the  solid  material  disappears ;  the  ash  remaining  differs 
from  the  original  fuel  in  almost  every  way;  a  flame 
attends  the  change ;  and  both  heat  and  light  are  given  off. 

These  observations  suggest  a  number  of  questions: 
What  is  the  cause  of  these  profound  changes  ?  How  does 
the  weight  of  the  ash  compare  with  that  of  the  original 
fuel  ?  What  becomes  of  the  material  that  disappears  ? 
What  is  the  flame  ?  What  is  the  nature  of  the  heat  and 
light?  Do  these  have  weight?  We  shall  learn  many 
facts  that  partly  answer  these  questions,  but  our  knowl- 
edge is  as  yet  so  limited  that  many  questions  will  remain 
but  partly  answered  at  the  end  of  our  study. 

No  change  in  weight  during  burning.  Our  ordinary 
experience  in  burning  coal  would  lead  us  to  suppose 
that  a  given  substance  loses  much  of  its  weight  during 
burning,  for  the  ashes  of  coal  certainly  weigh  much  less 


MATTER  AND  ENEKGY 


than  the  coal  itself.  Experiments  carried  out  with  great 
care  show  that  this  conclusion  is  too  hasty.  It  is  well 
known  that  coal  will  not  burn  unless  air  is  present,  and 
this  fact  requires  explanation  before  we  are  in  a  position 
to  draw  conclusions  about  the  loss  of  weight  during  com- 
bustion. If  air  is  used  up  when  the  coal  burns,  it  may  be 
that  other  invisible  gases  are  formed,  and  we  shall  have  to 
know  how  much  matter  the  coal  loses  in  this  way.  Some 
experiments  will  throw  much  light  upon 
this  subject. 

Experiment  1.  If  we  pour  a  little  clear 
limewater  into  a  wide-mouthed  bottle, 
insert  a  cork,  and  shake  the  bottle,  we 
see  little  or  no  change  in  the  liquid.  If, 
now,  we  hold  a  similar  empty  bottle 
mouth  downward  over  the  flame  of 
a  burning  candle  for  a  few  moments 
(Fig.  1),  moisture  will  be  seen  to  col- 
lect on  the  inside  of  the  cold  vessel, 

showing    that  water  vapor  is   formed 
,      .         ,          .  T(.  ,,  .  ln 

during   burning.     If  we   then    quickly 

pour  a  few  cubic  centimeters  of  clear 

limewater  into   the    bottle,    insert  the 

stopper,  and  shake  the  bottle  as  before,  we  notice  that  the 

clear  liquid  becomes  cloudy.    Some  invisible  gas  has  been 

formed  during  the  burning  of  the  candle,  has  risen  into 

the  bottle,  and  has  produced  a  change  in  the  clear  lime- 

water.    This  gas  is  called  carbon  dioxide. 

Experiment  2.  The  conclusion  we  have  reached  may  be 
confirmed  by  a  second  experiment.  A  candle  A  (Fig.  2) 
is  arranged  on  one  pan  of  a  large  balance  and  over  it  is 
suspended  a  lamp  chimney  or  cylinder  (i?)  open  at  both 
ends  and  closely  filled  with  pieces  of  soda  lime  (a  solid 


FIG.  1.  Collecting  the 
products  f  ormed  ?rom 

a  burning  candle 


AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


FIG.  2.  Demonstration  of  increase 
in  weight  during  burning 


substance  that  absorbs  both  water  vapor  and  carbon  diox- 
ide). The  candle  and  cylinder  are  then  balanced  by  the 
weights  C.  As  the  candle  burns,  the  pan  on  which  it  rests 
sinks,  showing  that  the  weight 
of  the  products  formed  by 
the  burning  of  the  candle  and 
absorbed  by  the  soda  lime 
is  greater  than  the  weight 
of  the  portion  of  the  candle 
burned. 

The  experiments  of  Lavoisier. 
Long  before  he  begins  the 
study  of  chemistry  every  stu- 
dent learns  that  the  air  con- 
tains an  invisible  gas  called 
oxygen.  Soon  after  the  dis- 
covery of  oxygen  as  a  con- 
stituent of  the  air  the  great  French  chemist,  Lavoisier 
(see  frontispiece),  made  some  experiments  that  enabled 
him  to  determine  the  changes  taking  place  when  a  sub- 
stance burns,  or,  as  it  is  often  expressed,  when  a  substance 
undergoes  combustion.  He  placed  some  pieces  of  solid 
materials  (tin,  iron,  phosphorus)  in  a  flask  full  of  air, 
closed  the  flask,  and  weighed  the  whole.  He  then  heated 
the  flask  and  its  contents  until  combustion  took  place,  not- 
ing that  not  all  of  the  solid  had  burned.  When  the  burning 
had  ceased  he  cooled  the  flask  and  weighed  the  whole  once 
more  and  found  that  there  had  been  no  change  in  weight 
as  the  result  of  combustion,  although  both  heat  and  light 
had  been  evolved,  and  heat  had  been  applied  to  the  flask. 
On  opening  the  flask,  air  rushed  in,  showing  that  the 
solid  in  burning  had  absorbed  the  air;  but  the  volume  of 
air  that  entered  was  much  less  than  the  volume  originally 


MATTER  AND  ENERGY  5 

present,  proving  that  not  all  of  the  air  had  been  absorbed 
during  combustion.  The  flask  was  again  weighed,  the 
increase  in  weight  being  the  weight  of  the  air  that  had 
entered  the  flask.  The  solids,  including  the  products  of 
combustion  and  the  unburned  materials,  were  then  weighed, 
and  the  weight  was  found  to  be  greater  than  before  com- 
bustion. Tlie  gain  in  weight  was  found  to  be  exactly  equal  to 
the  weight  of  the  air  used  in  the  burning.  By  other  experi- 
ments Lavoisier  was  able  to  show  that  when  a  substance 
burns  in  air,  not  all  the  constituents  of  the  air  are  used 
up,  but  only  the  oxygen  present. 

Results  of  the  experiments.  These  experiments  proved 
many  important  facts  —  among  them  that  the  act  of  com- 
bustion depends  upon  the  oxygen  of  the  air  as  well  as 
upon  the  material  undergoing  combustion ;  that  the  total 
weight  of  the  materials  concerned  is  not  changed  during 
combustion ;  and  that  the  light  and  heat  given  off  during 
combustion  occasion  no  loss  of  weight^  so  these  can  have 
no  weight. 

Matter  and  energy.  As  the  result  of  much  experience 
we  have  come  to  recognize  that  in  all  changes  similar  to 
combustion  there  are  two  fundamentally  different  things 
concerned :  (1)  the  material  that  possesses  weight  or  mass 
and  which  we  call  matter;  (2)  the  light  and  heat  given 
off,  which  have  no  weight  and  which  we  call  energy.  We 
shall  now  examine  these  two  somewhat  more  closely. 

Forms  of  energy.  We  sometimes  say  of  a  man  that  he 
is  full  of  energy,  meaning  that  he  has  a  great  capacity 
for  work.  Indeed  energy  is  often  defined  as  capacity  for 
work,  or  ability  to  do  work,  although  no  simple  definition 
as  yet  proposed  is  entirely  satisfactory.  We  recognize 
this  same  capacity  for  work  in  inanimate  things.  Steam, 
highly  compressed  in  a  boiler,  possesses  energy,  for  on 


6       AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

being  admitted  to  the  cylinder  of  a  locomotive  it  will 
push  back  the  piston  and  thus  move  the  train.  Energy 
is  present  in  electrical  power  lines,  for  through  the  neces- 
sary mechanical  devices  we  can  obtain  power,  heat,  and 
light  from  this  source.  A  moving  body  possesses  kinetic 
energy,  and  if  the  body  is  suddenly  stopped,  this  energy 
appears  as  heat. 

Conservation  of  energy.  The  experience  gained  in  a 
century  of  experimenting  has  convinced  scientists  that  it 
is  impossible  to  alter  the  quantity  of  energy  in  a  system 
of  bodies,  save  as  we  add  energy  from  without  or  allow 
it  to  escape  from  the  system,  and  this  generalization  is 
known  as  the  law  of  the  conservation  of  energy.  It  is  not 
difficult,  however,  to  transfer  energy  from  one  body  to 
another.  If  a  piece  of  hot  metal  is  dipped  into  water, 
the  metal  is  cooled  and  the  water  is  heated,  so  that  the 
metal  loses  energy  and  the  water  gains  it.  When  a  swing- 
ing bat  strikes  the  ball,  the  ball  gains  energy  while  the 
bat  loses  it.  It  is  evident,  therefore,  that  a  given  body 
does  not  possess  a  constant  quantity  of  energy  as  it  pos- 
sesses constant  mass. 

Transformation  of  energy.  Moreover,  energy  can  be 
freely  transformed  from  one  kind  into  another.  The 
heat  energy  of  burning  coal  can  be  changed  into  the 
kinetic  energy  of  the  locomotive.  The  kinetic  energy  of 
falling  water  can  be  transformed  into  electrical  energy, 
as  in  the 'power  plants  of  Niagara  Falls.  The  electrical 
energy  of  the  trolley  line  is  readily  converted  into  the 
kinetic  energy  of  the  moving  car.  In  all  such  transfor- 
mations a  definite  quantity  of  energy  of  one  kind  always 
gives  a  definite  quantity  of  another,  so  that  we  speak  of 
the  mechanical  equivalent  of  heat,  or  of  the  electrical 
equivalent  of  mechanical  energy. 


MATTER  AND  ENERGY  7 

The  diagram  (Fig.  3)  illustrates  a  few  familiar  transfor- 
mations of  energy.  The  heat  of  the  flame  A  is  converted 
into  mechanical  energy  in  the  heat  engine  B.  The  motion 
of  the  engine  is  communicated  to  the  small  dynamo  (7, 
where  it  is  converted  into  magnetic  and  electrical  energy. 
The  electrical  energy  is  changed  into  heat  and  light  in 
the  incandescent  lamp  Z>,  and  into  chemical  energy  (see 
following  paragraph)  by  the  decomposition  of  water  in  E. 


FIG.  3.   Diagram  illustrating  some  transformations  of  energy 

Chemical  energy.  A  body  may  possess  energy  due  to 
its  motion  or  to  its  position.  A  piece  of  coal,  however, 
possesses  energy  due  neither  to  its  motion  nor  to  its  position 
but  to  its  ability  to  undergo  combustion,  for  in  this 
process  both  heat  and  light  are  evolved.  Our  experience 
teaches  us  to  believe  that  this  heat  and  light  must  have 
come  from  some  other  form  of  energy  present  in  the  coal 
and  the  oxygen  which  unite  in  the  process  of  burning. 
This  form  of  energy  is  called  chemical  energy.  It  is  the 
form  possessed  by  substances  which  enables  them  to 
undergo  changes  similar  to  combustion,  and  it  is  the  form 
of  energy  in  which  the  chemist  is  especially  interested. 


8       AN  ELEMENTAEY  STUDY  OF  CHEMISTRY 


The  measurement  of  energy.  Since  changes  in  energy 
are  so  constantly  taking  place  all  about  us,  it  is  a  matter 
of  great  practical  importance  to  devise  units  for  the  meas- 
urement of  energy,  and  methods  for  making  the  measure- 
ment. In  general,  each  kind  of  energy  must  have  its  own 
units  of  measurement,  just  as  with  matter  we  have  centi- 
meters for  lengths,  liters  for  volumes,  and  grams  for 
weights.  In  some  of  its  forms  energy  is  very  difficult  to 
measure  directly,  and  neither  units 
nor  methods  for  the  direct  meas- 
urement of  chemical  energy  have 
as  yet  been  devised.  In  such  cases 
it  is  necessary  to  transform  the 
energy  into  a  form  more  conven- 
ient for  measurement.  In  the  case 
of  chemical  energy  it  is  changed 
into  heat  or  electrical  energy  for 
this  purpose. 

Measurement  of  heat.  A  quan- 
tity of  heat  energy  is  measured  by 
observing  to  what  extent  it  will 

change  the  temperature  of  a  given  mass  of  some  standard 
substance.  Water  has  been  chosen  as  the  standard,  and 
the  unit  of  heat  is  called  the  calorie  (designated  by  the 
abbreviation  cal.~).  It  is  denned  as  the  quantity  of  heat 
required  to  change  the  temperature  of  one  gram  of  water  one 
on  the  centigrade  scale. 


FIG.  4.   A  calorimeter 


The  actual  measurement  of  the  quantity  of  chemical  energy 
transformed  into  heat  in  any  definite  change  is  accomplished 
by  the  use  of  an  apparatus  called  the  calorimeter,  represented 
in  Fig.  4.  The  change  is  arranged  to  take  place  in  solution  in 
a  measured  volume  of  water  contained  in  a  thin-walled  metal 
vessel  A.  This  is  placed  within  a  double- walled  vessel  B, 


MATTER  AND  EKEHOY  9 

which  contains  water  at  the  temperature  of  the  room.  The 
thermometer  C  indicates  when  the  water  has  reached  this 
temperature.  This  water  is  to  prevent  the  influence  of  heat 
from  without,  and  as  an  added  precaution  the  vessel  is  covered 
with  a  thick  layer  of  nonconducting  felt.  The  heat  evolved  by 
the  change  raises  the  temperature  of  the  solution,  the  rise 
being  indicated  by  the  thermometers  D,  D.  During  the  change 
the  solution  is  stirred  by  the  stirrer  E.  If  the  weight  of  the 
water  is  (say)  2570  grams  and  the  rise  in  temperature  is  1.5° 
centigrade,  the  heat  evolved  is  2570  x  1.5  =  3855  cal. 

Matter  and  its  properties.  By  matter  we  mean  anything 
that  possesses  weight,  or  mass.  When  a  material  object 
is  at  rest  it  takes  an  expenditure  of  energy  to  set  it  in 
motion,  and  we  express  this  fact  by  saying  that  matter 
possesses  inertia.  Similarly,  when  a  body  is  in  motion  it 
requires  an  expenditure  of  energy  to  stop  it,  or  to  over- 
come its  inertia.  By  the  term  property  we  mean  a  mark 
or  characteristic  by  which  we  identify  a  given  thing.  Each 
form  of  matter  has  many  properties  peculiar  to  itself,  such 
as  its  physical  state  (solid,  liquid,  or  gas),  its  hardness, 
density,  luster,  and  shape ;  but  mass  and  inertia  are  the 
two  characteristics  or  properties  that  all  kinds  of  matter 
have  in  common. 

Law  of  conservation  of  matter.  Like  energy,  matter  may 
oftentimes  be  changed  from  one  form  into  another ;  thus  we 
have  seen  that  the  solid  candle  burns  away,  giving  rise  to 
invisible  gases.  The  solid  ice  is  easily  changed  into  liquid 
water  and  gaseous  steam.  The  question  naturally  arises 
whether  there  is  any  increase  or  decrease  in  the  weight 
of  the  matter  as  a  result  of  these  changes.  Much  careful 
experimenting  has  shown  that  there  is  not.  The  weight 
of  the  products  formed  in  any  change  in  matter  always 
equals  the  weight  of  the  substances  undergoing  change. 


10     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

We  may  alter  the  form  of  matter  but  not  its  mass.  This 
important  truth,  known  as  the  law  of  the  conservation  of 
matter,  may  be  stated  thus:  In  any  changes  to  which  we 
may  subject  a  given  quantity  of  matter,  the  mass  remains 
unchanged. 

Units  used  and  their  abbreviations.  In  the  chapters  to 
follow,  all  temperature  readings  given  will  be  those  of  the 
centigrade  scale,  unless  otherwise  designated.  In  referring 
to  other  units  of  measurements,  the  commonly  accepted 
abbreviations  will  be  employed,  such  as  g.  for  gram,  kg. 
for  kilogram,  cc.  for  cubic  centimeter,  1.  for  liter,  and  Ib. 
for  pound. 

EXERCISES 

1.  Give  instances  of  the  overlapping  of  different  sciences,  taken 
from  your  own  experiences  in  the  study  of  these  sciences. 

2.  Point  out  the  fundamental  difference  in  the  changes  involved 
in  (a)  the  crushing  of  a  stone  and  (ft)  the  burning  of  a  piece  of  coal. 

3.  How  do  the  ashes  of  coal  compare  in  weight  with  the  original 
coal?    Why  do  some  coals,  when  burned,  leave  more  ashes  than 
others  ? 

4.  How  do  you  account  for  the  fact  that  when  a  lamp  is  lighted, 
a  film  of  moisture  is  deposited  on  the  chimney  ?   Is  the  film  de- 
posited on  the  inside  or  the  outside  of  the  chimney  ?    Why  does  the 
film  soon  disappear  ? 

5.  Suggest  the  reason  why  moisture  often  collects  on  windows 
of  rooms  heated  by  gas  stoves  or  grates. 

6.  Narrate  some  of  the  important  incidents  in  the  life  of  Lavoi- 
sier (consult  encyclopedia). 

7.  Give  instances  illustrating   the   law  of  the  conservation  of 
energy,  and  the  law  of  the  conservation  of  matter. 

8.  The  energy  of  the  falling  water  at  Niagara  Falls  is  utilized 
in   propelling,  heating,  and  lighting   the  trolley  cars    at   Buffalo. 
Trace  the  changes  of  energy  involved. 

9.  What  is  the  source  of  the  energy  of  the  body? 


MATTER  AND  ENERGY  11 

10.  Why  does  the  body  become  warm  with  exercise? 

11.  Trace  the  energy  changes  between  the  coal  on  the  locomotive 
and  the  sound  of  the  whistle  that  one  hears. 

12.  In   Experiment  1,  why  does  the  water  vapor  change  into 
visible  moisture  on  the  sides  of  the  bottle? 

13.  How  many  forms  of  energy  can  you  name? 

14.  How  would  you  define  the  term  mechanical  equivalent  of  heat  1 

15.  The  fuel  value  of  a  coal  was  determined  by  burning  1  g.  of 
the  coal  in  a  calorimeter  containing  2500  g.  of  water.    The  heat 
liberated  raised  the  temperature  of  the  water  1.5°.    Calculate  the 
number  of  calories  of  heat  evolved  by  the  burning  of  the  coal. 

16.  A  certain  fuel  gives  out  4000  cal.  per  gram  when  burned. 
What  weight  of  it  would  be  required  to  heat  3  1.  of  water  from  room 
temperature  (18°)  to  the  boiling  point,  supposing  that  all  the  heat 
evolved  is  used  in  raising  the  temperature  of  the  water? 


CHAPTER  II 

VARIETIES  OF  MATTER:   COMPOUNDS,  ELEMENTS, 
MIXTURES 

Varieties  of  matter.  The  variety  of  forms  which  matter 
assumes  in  all  the  wonderful  transformations  of  nature  is 
almost  infinite,  and  these  forms  may  be  classified  in  a  great 
many  ways,  according  to  the  purpose  in  view.  The  inter- 
est of  the  chemist  centers  chiefly  in  the  composition  of 
substances  and  in  their  chemical  energy,  together  with 
the  changes  which  take  place  in  both  of  these.  From  this 
standpoint  he  finds  it  convenient  to  arrange  matter  in 
three  groups;  namely,  compounds,  elements,  and  mixtures. 

Illustrative  experiments.  The  distinction  between  these 
three  classes  can  be  explained  best  by  the  following 
illustrative  experiments. 

Experiment  1.  The  chief  properties  of  the  substances 
iron  and  sulfur  are  familiar  to  almost  everyone.  Iron 
filings  form  a  heavy  gray  powder,  strongly  attracted  by 
the  magnet.  When  treated  with  the  liquid  known  as  hy- 
drochloric acid,  the  iron  passes  into  solution  and  a  color- 
less gas  called  hydrogen  is  evolved,  considerable  heat 
being  liberated  in  the  process.  Sulfur  may  be.  obtained 
as  a  light  yellow  powder  not  attracted  by  a  magnet  nor 
dissolved  by  hydrochloric  acid.  It  is  readily  soluble  in 
the  liquid  known  as  carbon  disulfide,  however,  while  iron 
is  not  soluble,  and  when  the  solution  is  allowed  to  evap- 
orate, the  sulfur  is  deposited  in  the  form  of  yellow  crystals. 
12 


COMPOUNDS,  ELEMENTS,  MIXTURES 


13 


When  iron  filings  and  sulfur  are  thoroughly  ground 
together,  a  greenish-black  powder  is  obtained  which  in 
appearance  is  quite  different  from  either  of  them;  but 
when  we  apply  the  tests  which  we  have  found  to  char- 
acterize iron  and  sulfur,  it  is  found  that  in  many  respects 
the  powder  acts  like  these  two  taken  separately.  Hydro- 
chloric acid  still  dissolves  the  iron  and  evolves  hydrogen 
with  the  same  heat  as  before,  leaving  the  sulfur  unchanged. 
A  magnet  rubbed  through  the 
material  withdraws  the  iron  and 
leaves  the  sulfur.  Carbon  disul- 
fide  dissolves  the  sulfur  but  not 
the  iron.  The  sulfur  and  the 
iron  each  act  just  as  they  did  be- 
fore they  were  ground  together, 
and  with  the  same  energy. 

If,  now,  a  portion  of  the  pow- 
der is  placed  in  a  test  tube  and 

heated,  as  shown  in  Fig.  5,  it  FlG>  6<  Heating  a  mixture  of 
soon  begins  to  glow  at  the  point 
of  greatest  heat,  and  even  if  the 
flame  is  withdrawn,  the  glow  continues  to  spread  through- 
out the  entire  contents  of  the  test  tube,  and  a  great  deal 
of  heat  is  set  free  at  the  same  time.  When  the  product  is 
examined,  it  is  found  that  many  of  the  characteristics  of 
the  iron  and  sulfur  have  been  modified.  Carbon  disulfide 
no  longer  dissolves  sulfur  and  leaves  iron ;  a  magnet  has 
no  effect  upon  the  material ;  hydrochloric  acid  dissolves 
the  entire  product  and  evolves  a  gas  of  disagreeable  odor 
quite  different  from  hydrogen ;  and  the  heat  liberated  is 
quite  different  in  quantity  from  that  in  the  former  case. 
The  new  product  also  differs  from  the  iron  and  the  sulfur 
in  density,  color,  hardness,  solubility,  and  melting  point. 


iron  and  sulfur 


14     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  product  which  is  formed  by  the  union  of  iron  and 
sulfur  is  called  iron  sulfide. 

Experiment  2.  When  a  small  quantity  of  sugar  is 
heated  in  a  test  tube,  it  melts,  turns  brown  in  color, 
gives  off  vapors,  and  finally  dries  up  to  a  solid  black 
residue  which  may  be  identified  as  carbon.  By  collect- 
ing and  examining  the  vapors  it  is  possible  to  show  that 
they  are  largely  water.  To  make  this  transformation  com- 
plete it  is  necessary  to  apply  heat  throughout  the  entire 
process.  In  this  experiment  one  substance,  merely  by 
the  application  of  heat,  has  given  rise  to  two  others  of 
very  different  properties,  and  the  change  is  described  as 
a  decomposition. 

Chemical  reactions.  The  two  experiments  that  have 
just  been  described  are  very  different  in  many  ways,  but 
they  have  several  characteristics  in  common.  Most  of  the 
properties  of  the  materials  concerned  undergo  a  very  con- 
siderable change,  so  that  the  products  formed  are  different 
from  the  original  materials.  A  more  important  character- 
istic is  that  the  chemical  energy  of  the  materials  has  been 
changed.  The  action  of  iron  with  sulfur,  when  once 
started,  is  attended  by  the  evolution  of  a  great  deal  of 
heat,  and  this  is  at  the  expense  of  the  chemical  energy 
of  the  original  substances.  To  effect  the  decomposition 
of  sugar  it  is  necessary  to  supply  heat  during  the  entire 
period  of  decomposition,  and  this  heat  must  be  converted 
into  some  other  form  of  energy.  A  part  of  it  is  converted 
into  chemical  energy,  so  that  the  carbon  and  the  water 
taken  together  represent  more  energy  of  this  kind  than 
does  the  original  sugar.  Any  change  in  the  composition 
of  matter,  whether  union  or  decomposition,  that  involves 
a  change  in  the  chemical  energy  of  the  substances  con- 
cerned is  called  a  chemical  action  or  a  chemical  reaction. 


COMPOUNDS,  ELEMENTS,  MIXTURES  15 

Definition  of  chemical  compounds.  When  it  can  be 
shown  that  a  substance  is  composed  of  at  least  two 
different  materials,  and  that  its  chemical  energy  is  dif- 
ferent from  that  of  the  constituents  taken  separately,  it 
is  called  a  chemical  compound.  Thus,  we  judge  iron  sulfide 
to  be  a  chemical  compound,  for  it  contains  iron  and  sulfur, 
and  yet  differs  from  them  in  chemical  energy.  It  is  not 
always  a  simple  matter  to  determine  whether  a  given 
material  is  a  chemical  compound  or  not.  We  are  assisted 
in  our  decision  by  the  fact,  to  be  proved  in  a  later  chap- 
ter, that  the  percentage  composition  of  a  given  compound 
is  always  the  same.  Thus,  iron  sulfide  always  contains 
63.52  per  cent  iroft  and  36.48  per  cent  sulfur.  If  the 
material  can  be  obtained  in  pure  form  and  analyzed,  it  is 
usually  possible  to  decide  whether  or  not  it  is  a  compound 
by  the  constancy  of  its  composition. 

Chemical  affinity.  It  is  important  to  distinguish  clearly 
between  chemical  action  and  the  force  that  brings  about 
the  action.  This  force  is  called  chemical  affinity.  For 
example,  we  say  that  iron  and  sulfur  combine  because  of 
their  chemical  affinity.  We  can  form  little  idea,  as  yet, 
as  to  the  nature  of  this  force,  just  as  we  have  little  idea 
as  to  the  nature  of  the  force  of  gravitation.  In  both  cases 
we  merely  give  names  to  forces  which  we  must  believe 
to  be  acting. 

Conditions  affecting  chemical  action.  There  are  many 
conditions  which  may  either  promote  or  hinder  chemical 
action.  An  increase  hi  temperature  is  usually  favorable 
to  chemical  action,  as  was  seen  in  the  case  of  iron  and 
sulfur.  It  frequently  promotes  decomposition,  as  in  the 
case  of  sugar.  Other  forms  of  energy,  such  as  light, 
mechanical  pressure,  shock,  and  electrical  energy,  may 
also  facilitate  either  chemical  union  or  decomposition,  at 


16    AK  ELEMEHTABY  STUDY  OF  CHEMISTRY 

times  overcoming  obstacles  which  prevent  union,  in  other 
cases  overpowering  the  chemical  affinity  which  holds  a 
compound  together. 

Chemical  conduct  of  substances.  Substances  differ  very 
greatly  from  each  other  in  the  way  in  which  they  act 
toward  other  substances.  Thus,  a  substance  may  burn 
with  a  flame  when  heated  in  the  air,  or  it  may  combine 
with  another  substance  with  incandescence,  as  is  the  case 
of  iron  heated  with  sulfur.  It  may  decompose  when  heated, 
as  is  true  of  sugar,  or  when  subjected  to  the  action  of  the 
electric  current,  as  water  does.  All  such  peculiarities  are 
collectively  called  the  chemical  conduct  of  a  substance. 

Elements.  We  have  seen  that  sugar' can  be  decomposed 
into  two  different  substances,  namely,  water  and  carbon. 
\The  question  naturally  arises  whether  or  not  the  water 
and  carbon,  as  well  as  other  forms  of  matter,  such  as  iron 
and  sulfur,  can  likewise  be  decomposed  into  other  sub- 
stances.! To  determine  whether  or  not  any  given  sub- 
stance can  be  decomposed,  we  may  heat  the  substance,  as 
in  the  case  of  sugar;  or  we  may  employ  other  agencies. 
For  example,  experience  has  shown  that  in  many  cases 
decomposition  may  be  brought  about  by  the  electrical 
current  or  by  the  action  of  substances  possessed  of  great 
chemical  energy,  and  we  may  also  employ  these  methods. 
In  such  ways  chemists  have  succeeded  in  decomposing 
water  into  two  invisible  gases,  oxygen  and  hydrogen,  so 
that  water  must  be  regarded  as  a  compound. 

The  decomposition  of  water.  The  decomposition  of  water  may 
readily  be  observed  by  the  aid  of  an  apparatus  such  as  that 
represented  in  Fig.  6.  Two  test  tubes,  A  and  B,  are  filled 
with  water  and  inverted  in  a  vessel  half  filled  with  water  to 
which  a  little  sulfuric  acid  has  been  added.  A  piece  of  plat- 
inum foil,  C  and  D,  attached  to  a  wire  is  then  brought  under 


COMPOUNDS,  ELEMENTS,  MIXTURES          17 


the  end  of  each  tube.  When  these  wires  are  connected  with  a 
source  of  current  supplying  from  6  to  10  volts,  bubbles  of  gas 
will  be  seen  to  form  in  each  tube.  These  gases  may  be  shown 
to  have  different  properties ;  they  are  hydrogen  and  oxygen. 
The  reason  for  adding  the  sulfuric  acid 
will  be  discussed  later. 

On  the  other  hand,  carbon,  iron, 
and  sulfur  have  never  been  decom- 
posed, notwithstanding  the  many 
efforts  directed  to  this  end.  Sub- 
stances like  these  three,  which  have 
never  been  decomposed  into  two  or 
more  different  materials,  are  called 
elements,  or  elementary  substances.  It 
should  be  carefully  noted,  however, 
that  this  definition  does  not  suggest 
anything  as  to  the  real  nature  of  an 
element.  Neither  does  it  preclude 
the  possibility  that  one  element  may 
be  transformed  into  another  under 
some  condition  which  changes  its  chemical  energy.  In  the 
discussion  of  radium  it  will  be  shown  that  this  element 
is  slowly  changing  into  others. 

Number  of  the  elements.  While  many  thousands  of  com- 
pounds have  been  described,  the  number  of  the  elements 
at  present  known  is  comparatively  small.  A  complete  list 
is  given  in  the  table  on  the  inside  back  cover  of  this  book. 
Following  the  name  of  each  element  in  the  table  is  an  ab- 
breviation called  a  symbol,  by  which  the  element  is  desig- 
nated among  chemists.  The  symbol  is  usually  the  initial 
letter  of  the  name  of  the  element,  together  with  some 
other  characteristic  letter.  In  the  case  of  some  of  the 
elements  the  symbol  is  the  abbreviation  of  the  old  Latin 


FIG.  6.     The    decompo- 
sition of  water  into  oxy- 
gen and  hydrogen  by  the 
electric  current 


18     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

name,  as  is  true  of  iron  (ferrum),  gold  (aurum),  and 
mercury  {hydrargyrum).  The  significance  of  the  column 
of  numbers  will  be  made  clear  a  little  later. 

Physical  state  of  the  elements.  About  ten  of  the  ele- 
ments are  gases  at  ordinary  temperatures.  Two — mercury 
and  bromine- — are  liquids.  The  others  are  all  solids, 
though  their  melting  points  vary  through  wide  limits, 
from  caesium,  which  melts  at  26°,  to  elements  which  do 
not  melt  save  in  the  intense  heat  of  the  electric  furnace. 

Occurrence  of  the  elements.  Comparatively  few  of  the 
elements  occur  as  uncombined  substances  in  nature,  most 
of  them  being  found  in  the  form  of  chemical  compounds. 
When  an  element  does  occur  by  itself,  as  is  the  case  with 
gold,  we  say  that  it  occurs  in  the  free  state,  or  native ; 
when  it  is  combined  with  other  substances  in  the  form 
of  compounds,  we  say  that  it  occurs  in  the  combined  state, 
or  in  combination.  In  the  latter  case  there  is  usually 
little  about  the  compound  to  suggest  that  the  element  is 
present  in  it;  for  we  have  seen  that  elements  lose  their 
own  peculiar  properties  when  they  enter  into  combination 
with  other  elements.  From  its  appearance  it  would  never 
be  suspected  that  the  reddish,  earthy-looking  iron  ore 
contains  iron. 

Names  of  the  elements.  The  names  given  to  the  ele- 
ments have  been  selected  in  a  great  many  different  ways. 

(1)  Some  names  are  very  old  and  their  original  meaning 
is    obscure.      Such    names    are    iron,    gold,    and    copper. 

(2)  Many  names  indicate  some  striking  physical  property 
of  the  element.     The  name  bromine,  for  example,  is  de- 
rived from  a  Greek  word  meaning  "  stench,"  referring  to 
the  extremely  unpleasant  odor  of  the  substance.    (3)  Some 
names    indicate    the   chemical   conduct   of   the   elements. 
Thus,  nitrogen  means  "the  producer  of  niter,"   nitrogen 


COMPOUNDS,  ELEMENTS,  MIXTURES  19 

being  a  constituent  of  niter,  or  saltpeter ;  argon  means 
"  lazy,  or  inert,"  the  element  being  so  named  because  of  its 
inactivity.  (4)  Other  elements,  as  germanium  and  stron- 
tium, are  named  from  countries  or  localities.  (5)  Still 
others,  as  tantalum,  suggest  characters  in  mythology. 

Distribution  of  the  elements.  So  far  as  we  can  judge, 
these  elements  are  of  very  unequal  occurrence  in  nature. 
It  must  be  remembered,  however,  that  our  knowledge  of 
the  earth's  composition  is  confined  to  what  is  a  compara- 
tively thin  surface  shell,  not  exceeding  a  few  miles  in 
thickness.  The  table  below,  prepared  by  F.  W.  Clarke 
and  based  on  the  analysis  of  representative  rocks  and 
minerals,  gives  an  estimate  of  the  composition  of  this  solid 
shell.  It  will  be  seen  that  nine  of  the  elements  are  esti- 
mated to  constitute  98.4  per  cent  of  the  shell.  Some  of 
the  elements  are  of  such  rare  occurrence  that  only  a  few 
grams  have  ever  been  isolated. 

COMPOSITION  OF  THE  EARTH'S  CRUST 

Oxygen 47.07%  Potassium 2.45% 

Silicon 28.06%  Sodium       .    .    .'.    .  2.43% 

Aluminium     ....  7.90%  Magnesium     .     .  ' .'    .  2.40% 

Iron       ......  4.43%  Hydrogen  .     .     ...  0.22% 

Calcium 3.44%  Other  elements   .     .     .  1.60% 

Elements  essential  to  life.  A  careful  examination  of 
the  materials  present  in  living  organisms  shows  that  only 
a  few  are  of  vital  importance  to  us.  The  following 
table,  compiled  by  Sherman,  indicates  the  average  com- 
position of  the  human  body.  It  is  possible  that  other 
elements  have  an  importance  which  we  do  not  realize, 
but,  so  far  as  we  can  judge,  these  are  the  only  ones 
upon  which  living  organisms  are  dependent. 


20     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

AVERAGE  COMPOSITION  OF  THE  HUMAN  BODY 


Oxygen 

65.00% 

Phosphorus 

1.00% 

Magnesium 

0.05% 

Carbon  .     . 

18.00% 

Potassium    . 

0.35% 

Iron       .     . 

0.004% 

Hydrogen  . 

10.00% 

Sulfur      .     . 

0.25% 

Iodine    .     . 

traces 

Nitrogen    . 

3.00% 

Sodium    .     . 

0.15% 

Fluorine     . 

traces 

Calcium     . 

2.00% 

Chlorine 

0.15% 

Silicon  .     . 

traces 

Mixtures.  It  is  quite  possible  to  prepare,  either  from 
elements  or  from  compounds,  or  from  both,  a  body  which 
is  not  itself  a  compound  but  is  merely  a  mixture.  Ordi- 
nary concrete  is  such  a  material,  for  in  a  broken  piece 
it  is  easy  to  identify  the  crushed  stone,  the  sand,  and  the 
cement  which  compose  it.  Granite  is  a  sort  of  natural 
concrete,  in  which  two  very  different-looking  crystalline 
materials,  mica  and  feldspar,  are  bound  together  by  a 
glassy  substance  called  silica.  Iron  and  sulfur  when 
rubbed  together  form  a  material  more  closely  resembling 
a  compound,  in  that  it  is  apparently  of  even  quality 
throughout,  or  is  homogeneous.  An  examination  under 
the  microscope  shows  that  this  is  not  really  so,  for  the 
particles  of  iron  and  sulfur  can  still  be  seen  lying  side 
by  side  unchanged. 

Alchemy.  In  olden  times  it  was  thought  that  some  way  could 
be  found  to  change  one  element  into  another,  and  a  great  many 
efforts  were  made  to  accomplish  this  transformation.  Most 
of  these  efforts  were  directed  toward  changing  the  commoner 
metals  into  gold,  and  many  fanciful  ways  for  doing  this., were 
described.  The  chemists  of  that  time  were  called  alchemic*. 
and  the  art  which  they  practiced  was  called  alchemy.  The 
alchemists  gradually  became  convinced  that  the  only  way 
common  metals  could  be  changed  into  gold  was  by  the  won- 
derful power  of  a  magic  substance  which  they  called  the  phi- 
losophers' stone,  which  would  accomplish  this  transformation 
by  its  mere  touch  and  would  in  addition  give  perpetual  youth 


COMPOUNDS,  ELEMENTS,  MIXTURES          21 

to  its  fortunate  possessor.  No  one  has  ever  found  such  a 
stone,  and  no  one  has  succeeded  in  changing  one  metal  into 
another. 

One  of  the  most  brilliant  discoveries  of  recent  years  seems 
in  one  sense  to  come  near  realizing  the  dreams  of  the  alche- 
mists, for  it  has  been  shown  .that  one  or  two  undoubted  ele- 
ments are  very  slowly  changing  into  others.  But  no  one  has 
been  successful  in  either  hastening  or  retarding  the  change. 


FIG.  7.   An  alchemist's  laboratory,  according  to  the  painting  of  Teniers 

The  domain  of  chemistry.  With  the  general  character- 
istics^of  matter  and  energy  before  us,  it  is  now  possible 
for  us  to  form  some  idea  of  the  topics  that  are  of  funda- 
mental interest  to  the  chemist.  In  the  first  place,  the 
chemist  is  concerned  with  the  composition  of  matter.  For 
example,  such  questions  as  whether  a  given  substance 
is  an  element  or  a  compound ;  how  it  may  be  obtained 
in  pure  form  and  what  its  properties  and  uses  are ; 


22     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

whether  a  certain  rock  contains  iron  or  gold;  whether  a 
given  sample  of  drinking  water  contains  anything  that  is 
injurious  to  health,  —  these  are  questions  for  the  chemist 
to  decide.  Moreover,  the  chemist  is  equally  interested  in 
the  changes  both  in  form  and  energy  which  any  given 
substance  undergoes.  It  is  of  fundamental  importance 
that  he  should  know  what  compounds  any  given  element 
will  form ;  what  the  nature  of  these  compounds  is ;  which 
of  two  given  samples  of  coal  evolves  the  greater  amount 
of  heat  on  burning ;  and  what  the  function  of  the  various 
classes  of  foods  is  and  what  -  mixture  of  these  is  best 
adapted  to  preserve  health  under  different  physical  con- 
ditions. It  is  with  questions  of  this  kind  that  we  shall 
be  concerned  in  the  study  of  chemistry. 


1.  What  means  of    decomposing  a  compound  can  you  suggest 
besides  heating? 

2.  Define  a  compound ;  an  element.    Which  group  is  the  more 
numerous  ? 

3.  What  is  meant  by  the  earth's  crust? 

4.  Does  the  fact  that  a  substance  undergoes  no  change  on  heat- 
ing show  it  to  be  an  element  ? 

5.  Read  over  the  list  of  elements.    What  ones  do  you  know  to 
occur  native? 

6.  Aluminium  is  much  more  abundant  than  iron  (see  Clarke's 
table).     How   do   you   account   for   the    much    greater   cheapness 
of  iron? 

7.  Consult  the   dictionary   for  the    derivation  and    significance 
of  the  names  of   the    following   elements :  phosphorus,  hydrogen, 
germanium,  columbium,  chlorine,  argon,  copper,  selenium,  thorium, 
iodine. 

8.  Give  examples  of  chemical  action  caused  through  the  agency 
of  heat;  of  light;  of  electricity. 


COMPOUNDS,  ELEMENTS,  MIXTURES          23 

9.  Calculate  the  approximate  weights  of  the  principal  elements 
present  in  your  body. 

10.  How  do  you  account  for  the  fact  that  the  mound  builders 
used  vessels  of  copper  rather  than  of  iron? 

11.  Give  reasons  why  gold  is  more  costly  than  iron. 

12.  Give  reasons  why  iron  does  not  occur  native  to  any  extent. 

13.  What  is  a  characteristic  property  of  all  the  elements  that 
occur  native? 

14.  What  is  meant  by  the  physical  state  of  an  element  ? 

15.  Would  the  fact  that  iron  is  present  in  the  human  body  in 
such  very  small  amounts  suggest  that  we  could  get  along  without  it  ? 

16.  Why  is  the  list  of  elements  subject  to  change  from  year  to 
year? 

17.  What  can  you  add  to  what  has  been  said  about  the  alchemists  ? 
(Consult  Encyclopedia.) 


CHAPTER  III 


OXYGEN 

Historical.  The  discovery  of  oxygen  is  attributed  to  the 
English  chemist  Priestley,  who  in  1774  obtained  the  ele- 
ment by  heating  a  compound  of  mercury  and  oxygen  now 

known  as  red  oxide  of  mer- 
cury. It  is  probable,  however, 
that  other  investigators,  es- 
pecially the  Swedish  chem- 
ist Scheele,  had  obtained  it 
at  an  earlier  date  but  failed 
to  attract  attention  to  their 
discovery.  The  name  oxygen 
signifies  "  acid  former."  It 
was  given  to  the  element  by 
the  French  chemist  Lavoisier, 
since  he  believed  that  all  acids 
owe  their  characteristic  prop- 
erties to  the  presence  of  oxy- 
gen. This  idea  we  now  know 
to  be  incorrect. 

Occurrence.  Oxygen  is  by 
far  the  most  abundant  of  all 
the  elements.  It  occurs  both  in  the  free  state  and  in  com- 
bination. In  the  free  state  it  is  found  in  the  air,  100  vol- 
umes of  dry  air  containing  about  21  volumes  of  oxygen. 
In  the  combined  state  it  forms  88.81  per  cent  of  water 
and  nearly  one  half  of  the  rocks  composing  the  earth's 
24 


FIG.  8.   Joseph  Priestley 
(1733-1804) 

The  discoverer  of  oxygen 


OXYGEN  25 

Crust.  It  is  also  an  important  constituent  of  the  com- 
pounds that  compose  plant  and  animal  tissues ;  for  ex- 
ample, about  two  thirds,  by  weight,  of  the  human  body 
is  oxygen. 

Preparation.  While  oxygen  is  very  abundant,  it  does 
not  occur  in  nature  in  pure  condition.  To  obtain  pure 
oxygen  we  must  either  liberate  it  from  some  compound 
or  separate  it  from  the"  gases  with  which  it  is  mixed  in 
the  air.  The  most  important  of  the  methods  for  preparing 
the  pure  element  are  the  following: 

1.  Preparation  from  water.    Water  is  a  compound,  con- 
sisting of  88.81  per  cent  oxygen  and  11.19  per  cent  hy- 
drogen.   It  is  easily  separated  into  these  constituents  by 
passing  an  electric  current  through  it,  as  has  been  already 
explained  (pp.  16,  17). 

2.  Preparation   by  heating  certain   compounds   of  oxygen. 
Some  of  the   compounds   of   oxygen,  when   heated,   give 
off  at  least  a  portion  of  their  oxygen.    For  example,  mer- 
curic oxide,  a  solid  compound  containing   7.39  per  cent 
of  oxygen  and  92.61  per  cent  of  mercury,  is  decomposed 
into  its  elements  by  heating  it.    The  change  may  be  repre- 
sented in  the  following  way,  in  which  the  names  of  the 
elements  composing  the  compound  are  inclosed  in  brackets 
just  beneath  the  names  of  the  compounds: 

mercuric  oxide >-  mercury  +  oxygen 

[mercury"! 
oxygen  J 

The  compound  best  adapted  for  the  preparation  of 
oxygen  by  this  method  is  potassium  chlorate.  This  com- 
pound is  a  white  solid  containing  39.16  per  cent  of 
oxygen,  31.90  per  cent  of  potassium,  and  28.94  per  cent 
of  chlorine.  When  heated  it  undergoes  a  series  of  changes 
in  which  all  the  oxygen  is  finally  set  free,  leaving  a  white 


26     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


solid   compound   of  potassium  and  chlorine  called  potas- 
sium chloride.   The  change  may  be  represented  as  follows : 

potassium  chlorate >•  potassium  chloride  +  oxygen 


[potassium 
chlorine 
oxygen 


[potassium! 
chlorine     J 


The  evolution  of  the  gas  becomes  marked  at  about 
400°.  It  is  a  remarkable  fact  that  the  rate  at  which  the 
oxygen  is  evolved  at  any  given  temperature  is  greatly 
increased  by  the  presence  of  small 
quantities  of  certain  substances,  not- 
ably manganese  dioxide.  By  mixing 
such  a  substance  with  the  chlorate  it 
is  possible,  therefore,  to  obtain  the 
oxygen  rapidly 
at  a  lower  tem- 
perature than 
would  otherwise 
have  to  be  em- 
ployed. As  to 


FIG.  9.    Preparation  of  oxygen  from  potassium 
chlorate,  and  method  of  collecting  the  gas 


the  way  in  which 
the  manganese 
dioxide  promotes  the  decomposition,  it  may  be  said  at  once 
that  we  do  not  know.  Apparently  it  undergoes  no  change 
during  the  reaction.  Certainly  it  contributes  no  oxygen, 
for  the  weight  of  oxygen  obtained  is  always  39.16  per  cent 
of  the  weight  of  the  chlorate  used,  irrespective  of  the  pres- 
ence of  manganese  dioxide.  This  is  but  one  example  of 
many  in  which  the  rate  of  change  is  influenced  by  an 
apparently  inactive  substance.  Such  materials  are  called 
catalytic  agents,  or  catalyzers,  and  we  shall  meet  with  them 
frequently  in  subsequent  pages,  since  a  great  many  chemical 
processes  depend  upon  suitable  catalyzers  for  their  success. 


OXYGEN  27 

Laboratory  preparation  of  oxygen.  The  preparation  of  oxygen 
from  potassium  chlorate  as  commonly  carried  out  in  the  labo- 
ratory is  as  follows : 

The  potassium  chlorate,  mixed  with  about  one  fourth  of  its 
weight  of  manganese  dioxide,  is  placed  in  a  suitable  vessel, 
such  as  a  glass  flask,  which  is  provided  with  a  stopper  and  glass 
tube,  as  shown  in  A  (Fig.  9).  Upon  applying  a  gentle  heat, 
oxygen  is  evolved  and  passes  out  through  the  tube  B.  It  is 
evident  that  the  oxygen  which  first  escapes  is  mixed  with  the 
air  contained  in  the  flask.  In  a  short  time,  as  the  evolution  of 
oxygen  continues,  all  this  air  is 
displaced,  and  the  pure  oxygen 
may  then  be  collected  by  bringing 
the  end  of  the  delivery  tube  under 
the  mouth  of  a  glass  cylinder  C, 
which  has  been  filled  with  water 
and  inverted  in  a  trough  of  water 
D,  as  shown  in  the  figure.  The  gas 
rises  in  the  cylinder  and  displaces  FIG.  10.  Copper  retort  for  mak- 
the  water.  In  preparing  larger  ing  oxygen 

quantities    of    oxygen    a   copper 

retort  (Fig.  10)  having  a  capacity  of  from  500  to  1000  cc.  may 
be  used  to  advantage  in  place  of  the  more  fragile  glass  flask. 

3.  Preparation  from  sodium  peroxide.  This  compound  is 
a  white  solid  containing  41.02  per  cent  of  oxygen.  When 
water  is  brought  into  contact  with  it,  the  two  react  in 
such  a  way  as  to  liberate  a  portion  of  the  oxygen.  This 
reaction  may  be  expressed  as  follows : 

sodium  peroxide  +  water >•  sodium  hydroxide  4-  oxygen 

[sodium!  [hydrogen!  ["sodium     "I 

LoxygenJ  [oxygen     J  hydrogen 

Loxygen     J 

While  this  is  an  expensive  method,  it  is  often  used  because 
of  its  simplicity.  It  is  only  necessary  to  bring  the  two  com- 
pounds together  in  order  to  obtain  the  gas. 


28     AN  ELEMENTABY  STUDY  OF  CHEMISTEY 

4.  Preparation  from  air.  Since  air  contains  such  a  large 
percentage  of  free  oxygen,  one  would  naturally  expect 
methods  to  be  devised  for  obtaining  oxygen  from  this 
source.  The  problem  is  not  so  simple  as  it  may  seem, 
for  there  are  other  gases  in  the  air,  and  the  separation 
of  a  gas  in  pure  condition  from  a  mixture  of  gases  is 
always  difficult. 

To  accomplish  the  separation,  the  air  is  first  subjected 
to  the  combined  effects  of  pressure  and  low  temperature. 
In  this  way  it  is  possible  to  obtain  it  in  the  form  of  a 
liquid  which  is  essentially  a  mixture  of  oxygen  and  nitro- 
gen in  the  liquid  state.  When  this  liquid  is  allowed  to 
stand  under  ordinary  pressure,  it  boils  rapidly  and  its 
temperature  falls  to  a  very  low  point.  Since  nitrogen  has 
the  lower  boiling  point  (—195.7°),  it  tends  to  boil  away 
first,  and  is  gradually  followed  by  the  oxygen  (boiling 
point, —182.9°),  which  may  be  collected  separately. 

Laboratory  and  commercial  methods  of  preparation.  We 
have  seen  that  a  number  of  different  methods  may  be 
used  for  the  preparation  of  oxygen.  We  shall  find  as 
we  proceed  that  this  is  true  in  reference  to  most  of  the 
other  elements.  Some  of  these  methods  are  expensive, 
while  others  necessitate  the  use  of  complicated  apparatus 
or  costly  machinery.  For  the  purpose  of  laboratory  experi- 
ments, in  which  relatively  small  quantities  are  desired,  the 
choice  of  the  method  will  naturally  be  decided  by  con- 
venience and  simplicity  of  apparatus,  while  in  the  prepara- 
tion on  a  commercial  scale  economy  will  determine  the 
method.  In  the  case  of  oxygen  the  method  of  preparation 
from  potassium  chlorate  has  proved  itself  the  most  suit- 
able for  laboratory  purposes.  For  commercial  purposes 
oxygen  is  obtained  either  from  water  (method  1)  or 
from  air  (method  4). 


OXYGEN  29 

Properties  of  oxygen.  Oxygen  is  a  colorless,  odorless, 
tasteless  gas,  slightly  heavier  than  air.  One  liter  of  it, 
measured  at  a  temperature  of  0°  and  under  a  pressure  of 
1  atmosphere,  weighs  1.4290  g.,  while  under  similar  con- 
ditions 1  1.  of  air  weighs  1.2928  g.  It  is  but  slightly 
soluble  in  water,  100  volumes  of  water  at  0°  and  under 
ordinary  atmospheric  pressure  dissolving  about  4  volumes 
of  the  gas.  Oxygen,  like  other  gases,  may  be  liquefied 
by  applying  very  great  pressure  to  the  highly  cooled  gas. 


FIG.  11.   Oxygen  stored  in  steel  cylinders 

When  the  pressure  is  removed,  the  liquid  oxygen  passes 
again  into  the  gaseous  state,  since  its  boiling  point  under 
ordinary  atmospheric  pressure  is  — 182.9°.  By  reducing 
the  temperature  still  lower  oxygen  is  obtained  in  the 
form  of  a  snowlike  solid  which  melts  at  —  235°.  For 
purposes  of  transportation  the  gas  is  pumped  under  great 
pressure  into  strong  steel  cylinders  (Fig.  11),  and  it  may 
be  purchased  in  this  form. 

Chemical  conduct.  At  ordinary  temperatures  oxygen  is 
not  very  active ;  most  substances  either  are  not  affected 
by  it  or  are  affected  so  slowly  that  the  action  escapes  notice. 
At  higher  temperatures,  however,  oxygen  is  very  active 


30     AN  ELEMENTARY  STUDY  OF  CHEMISTEY 


and  unites  directly  with  most  of  the  elements.  This  may  be 
shown  by  heating  various  substances  until  they  are  just 
ignited  in  air,  and  then  bringing  them  into  vessels  con- 
taining oxygen,  when  they  burn  with  greatly  increased 
brilliancy.  Thus,  a  glowing  splint  introduced  into  a  jar 
of  oxygen  bursts  into  flame.  Sulfur  burns  in  air  with  a 
very  weak  flame  and  a  feeble  light ;  in  oxygen  the  flame 
is  increased  in  size  and  brightness  (Fig.  12).  Substances 
which  burn  readily  in  the  air,  such  as 
phosphorus,  burn  in  oxygen  with  daz- 
zling brilliancy.  Many  substances  which 
burn  in  the  air  with  great  difficulty, 
such  as  iron,  burn  readily  in  oxygen. 
The  nature  of  the  action  of  oxygen 
upon  substances  ;  oxidation  ;  oxidizing 
agent.  It  is  possible  to  show  by  experi- 
ment that  the  action  of  oxygen  upon 
another  element  consists  in  the  union  of 
the  two  elements  to  form  a  compound. 
Thus,  when  sulfur  burns  in  oxygen, 
both  sulfur  and  oxygen  disappear  as 
such,  and  in  their  place  we  find  a  gas- 
eous compound  composed  of  the  two  elements.  Likewise, 
*when  phosphorus  or  iron  or  carbon  burn  in  oxygen,  com- 
pounds of  these  elements  with  oxygen  are  formed.  Many 
compounds  as  well  as  elements  burn  readily  both  in  air 
and  in  oxygen;  among  these  are  coal,  wood,  oil,  and  gas. 
In  the  majority  of  such  cases  the  compound  is  completely 
decomposed  and  each  of  its  constituent  elements  com- 
bines with  oxygen.  Thus,  most  oils  are  made  up  of  carbon 
and  hydrogen,  and  when  the  oil  burns  it  is  converted 
into  a  compound  of  carbon  and  oxygen  (carbon  dioxide) 
and  a  compound  of  hydrogen  and  oxygen  (water).  Less 


FIG.  12.  Burning  sul- 
fur in  oxygen 


OXYGEN  31 

frequently  the  compound  undergoes  no  decomposition  but 
merely  as  a  whole  combines  with  oxygen. 

The  general  term  oxidation  is  applied  to  all  such 
processes  as  those  described  above,  in  which  any  sub- 
stance or  its  constituent  parts  combines  with  oxygen. 
Thus,  we  speak  of  the  oxidation  of  phosphorus  or  sulfur 
by  the  air  or  by  pure  oxygen,  and  we  say  that  these 
elements  readily  undergo  oxidation.  The  material  which 
supplies  the  oxygen  is  called  the  oxidizing  agent.  In  the 
examples  just  mentioned  the  air  or  pure  oxygen  is  the 
oxidizing  agent,  but  in  many  cases  the  oxygen  is  sup- 
plied by  some  compound  such  as  potassium  chlorate  or 
sodium  peroxide. 

Oxides  ;  products  of  oxidation.  When  any  element  com- 
bines with  oxygen,  the  resulting  compound  is  known  as 
an  oxide  of  that  element.  Thus,  the  compound  formed  by 
the  union  of  sulfur  with  oxygen  is  known  as  an  oxide 
of  sulfur.  Likewise,  when  phosphorus  or  iron  or  carbon 
combine  with  oxygen,  the  resulting  compounds  formed 
are  oxides.  The  particular  oxide  or  oxides  formed  in  the 
oxidation  of  any  substance  are  known  in  general  as  the 
products  of  oxidation  of  that  substance. 

Oxides  of  nearly  all  the  elements  have  been  prepared, 
and  they  constitute  an  important  class  of  compounds.  Some 
of  these  oxides  are  invisible  gases,  as  is  true  of  the  oxides 
of  sulfur  and  of  carbon.  In  a  few  cases  the  oxide  is  a 
liquid,  the  most  familiar  example  being  water,  which  is 
an  oxide  of  hydrogen.  In  the  great  majority  of  cases, 
however,  the  oxides  are  solids,  which  is  true  of  those  of 
iron  and  phosphorus.  It  is  easy  to  understand,  therefore, 
why  such  elements  as  sulfur  and  carbon  completely  vanish 
on  burning,  leaving  no  ash,  while  other  elements,  such  as 
iron  and  phosphorus,  leave  a  solid  residue. 


32     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Combustion.  Sometimes  oxidation  takes  place  so  slowly 
that  no  light  is  seen  and,  unless  careful  measurements 
are  made,  no  heat  is  noticed.  The  decay  of  vegetable 
matter,  such  as  wood  and  leaves,  is  an  example  of  this 
slow  oxidation.  In  other  cases,  as  with  burning  phos- 
phorus or  iron,  light  is  given  off  either  as  a  flame  or  as 
a  glow  called  incandescence.  In  such  cases  the  substance 
is  said  to  undergo  combustion.  In  its  broad  sense  the 
term  combustion  is  applied  to  any  chemical  reaction  in 
which  light  is  evolved.  The  most  familiar  examples  of 
combustion  are  those  in  which  substances  burn  in  air 
or  oxygen,  and  which  are  therefore  also  oxidations.  Ordi- 
narily, when  we  speak  of  a  combustible  substance  we 
mean  one  that  will  burn  in  air  or  oxygen. 

Heat  of  oxidation  and  combustion.  Evidently  a  given 
substance  may  either  undergo  a  slow  oxidation  or  it  may 
undergo  combustion.  Thus,  a  piece  of  phosphorus,  ex- 
posed to  the  air  in  a  cold  room,  slowly  wastes  away 
until  it  has  all  disappeared  into  smoke  consisting  of  an 
oxide  of  phosphorus ;  but  if  it  is  touched  with  a  lighted 
match,  it  takes  fire  and  burns  very  rapidly,  giving  out 
much  heat  in  its  combustion.  The  product  is  the  same 
in  both  cases  ;  namely,  an  oxide  of  phosphorus.  Apparently 
the  difference  lies  in  the  amount  of  heat  given  off,  but 
very  accurate  experiments  demonstrate  that  this,  too,  is 
exactly  the  same.  In  the  one  case  the  action  is  so  slow 
that  the  heat  is  conducted  away  as  fast  as  it  is  liberated, 
and  so  it  escapes  notice;  in  the  other  it  is  given  off  so 
rapidly  as  to  be  very  striking.  A  similar  relation  has 
been  found  to  hold  true  in  all  cases  of  combustion.  The 
heat  given  off  when  a  definite  weight  of  a  substance  under- 
goes oxidation  is  exactly  the  same  whether  the  action  is 
fast  or  slow,  provided  the  same  compound  is  formed. 


I 


OXYGEN          .  33 

Spontaneous  combustion.  It  has  been  found  that  the 
rate  at  which  oxidation  goes  on  is  greatly  increased  tiy 
raising  the  temperature  of  the  material  undergoing  oxida- 
tion. Consequently,  if  the  conditions  surrounding  oxidation 
are  such  that  the  heat  given  off  cannot  escape,  the  tem- 
perature will  steadily  rise,  and  because  of  this  the  rate  of 
oxidation  will  increase.  The  increased  heat  thus  set  free 
will  still  further  raise  the  temperature,  until  the  oxidation 
passes  into  active  combustion,  the  point  at  which  this 
occurs  being  called  the  kindling  temperature.  Materials 
taking  fire  in  this  way  are  said  to  undergo  spontaneous 
combustion.  It  will  be  seen  that  the  essential  conditions 
are  (1)  an  existing  slow  oxidation  and  (2)  good  heat 
insulation.  Linseed  oil,  used  in  paints,  undergoes  rather 
rapid  oxidation  in  air,  and  oily  rags  left  by  painters  not 
infrequently  occasion  disastrous  fires.  Fine  coal  in  the  cen- 
ter of  a  heap  or  in  the  closed  hold  of  a  vessel  sometimes 
takes  fire.  Almost  any  finely  divided  combustible  ma- 
terial, such  as  sawdust  or  flour,  is  dangerous  when  stored 
in  a  warm  dry  place.  Sometimes  the  heat  of  fermentation, 
which  is  a  kind  of  oxidation,  will  start  a  fire  in  a  haystack 
or  barn  if  the  hay  is  not  well  dried  before  being  stored. 

Importance  of  oxygen.  The  great  importance  of  oxygen 
in  nature  is  evident  from  the  facts  which  have  already 
been  presented  in  this  chapter.  It  is  a  constituent  of  the 
great  majority  of  the  compounds  which  collectively  con- 
stitute the  solid  earth,  the  living  creatures  upon  it,  and 
the  water  which  covers  so  much  of  its  surface,  while  the 
atmosphere  is  a  great  reservoir  from  which  a  supply  of 
the  free  element  can  be  drawn  at  any  time. 

Free  oxygen  is  essential  to  the  life  of  all  organisms, 
with  the  exception  of  some  of  the  lowest  forms.  Aquatic 
animals  obtain  the  necessary  oxygen  from  the  air  dissolved 


34     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

in  the  water  in  which  they  live.  Free  oxygen  also  plays 
a  prominent  part  in  the  decomposition  of  refuse  organic 
matter,  much  of  it  being  oxidized  into  harmless  gases 
(Fig.  13).  It  is  noteworthy,  however,  that  the  oxidation 
of  such  matter  takes  place  only  in  the  presence  of  certain 
minute  forms  of  living  organisms  known  as  bacteria. 

Free  oxygen  is  also  utilized  in  a  great  variety  of  indus- 
trial processes,  but  for  most  of  these,  air  answers  every 


FIG.  13.    Sewage-disposal  plant  at  Columbus,  Ohio,  in  which  the  sewage 
is  sprayed  into  the  air  to  secure  its  oxidation 

purpose,  since  the  nitrogen  which  it  contains  does  not 
seriously  interfere.  Pure  oxygen  finds  application  in  quite 
a  variety  of  scientific  experiments,  in  the  production  of 
very  high  temperatures,  and  in  the  treatment  of  certain 
diseases  in  which  the  patient  is  unable  to  inhale  sufficient 
air  to  supply  the  necessary  quantity  of  oxygen. 

Phlogiston  theory  of  combustion.  The  French  chemist  Lavoi- 
sier (1743-1794),  who  gave  to  oxygen  its  name,  was  the  first 
to  show  that  combustion  is  due  to  union  with  oxygen.  Pre- 
vious to  his  time  combustion  was  supposed  to  be  due  to  the 


OXYGEN  35 

presence  of  a  substance  or  principle  called  phlogiston.  One  sub- 
stance was  thought  to  be  more  combustible  than  another  be- 
cause it  contained  more  phlogiston.  Coal,  for  example,  was 
thought  to  be  very  rich  in  phlogiston.  The  ashes  left  after 
combustion  would  not  burn  because  all  the  phlogiston  had 
escaped.  If  the  phlogiston  could  be  restored  in  any  way,  the 
substance  would  then  become  combustible  again.  Although 
this  view  seems  absurd  to  us  in  the  light  of  our  present  knowl- 
edge, it  formerly  had  general  acceptance.  The  discovery  of 
oxygen  led  Lavoisier  to  investigate  the  subject,  and  through 
his  experiments  he  arrived  at  the  true  explanation  of  combus- 
tion. The  discovery  of  oxygen,  together  with  the  part  it  plays 
in  combustion,  is  justly  regarded  as  one  of  the  most  important 
discoveries  in  the  history  of  chemistry.  It  marked  the  dawn 
of  a  new  period  in  the  growth  of  the  science. 

The  definiteness  of  chemical  processes.  Throughout  this 
chapter  attention  has  been  repeatedly  directed  to  the  fact 
that  chemical  processes  involve  definite  weights  of  matter. 
For  example,  the  composition  of  a  number  of  compounds 
has  been  expressed  in  exact  percentages,  since  experiment 
has  shown  that  these  always  have  precisely  the  composi- 
tion stated,  irrespective  of  the  source  from  which  they  are 
obtained  or  the  method  by  which  they  are  prepared.  After 
extensive  investigation  of  a  very  large  number  of  com- 
pounds, chemists  have  concluded  that  this  constancy  of 
composition  is  a  characteristic  of  every  true  compound, 
and  a  statement  of  this  characteristic  is  commonly  called 
the  law  of  definite  composition. 

In  like  manner,  the  chemical  changes  which  compounds 
undergo  are  always  perfectly  definite  under  stated  condi- 
tions. Thus,  when  potassium  chlorate  is  heated,  for  every 
100  g.  decomposed  there  result  39.16  g.  of  oxygen  and 
60.84  g.  of  potassium  chloride.  When  iron  burns  in  oxy- 
gen, 100  g.  of  iron  combines  with  38.20  g.  of  oxygen  to 


36     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

form  138.20  g.  of  oxide  of  iron.  If  less  than  38.20  g.  of 
oxygen  is  present,  then  a  corresponding  amount  of  iron 
will  remain  unchanged.  On  the  other  hand,  if  more  than 
38.20  g.  of  oxygen  is  present,  then  all  the  iron  will  be 
changed  into  the  oxide,  and  the  excess  of  oxygen  will 
remain  unaltered.  The  actual  experiments  which  justify 
these  conclusions  will  come  before  us  from  time  to  time 
as  we  proceed. 


In  all  the  problems  in  this  text  which  involve  volumes  of  gases,  it  is  under- 
stood that,  unless  otherwise  designated,  the  volumes  referred  to  are  those 
which  the  gas  will  occupy  at  a  temperature  of  0°  and  under  a  pressure 
of  1  atmosphere. 

1.  In  Fig.  9  why  does  the  water  stay  in  the  inverted  cylinder? 
Why  does  the  oxygen  displace  it  ?  When  a  little  oxygen  has  entered, 
why  does  not  all  the  water  run  out? 

2.  Suggest  a  method  for  collecting  a  gas  that  is  soluble  in  water. 

3.  Why  is  it  that  the  discovery  of  oxygen  is  ordinarily  attrih- 
iited  to  Priestley,  although  others  had  obtained  it  before  him  ? 

4.  Report  brief  accounts  of  the  lives  of  Priestley,  Scheele,  and 
Lavoisier  (consult  encyclopedia). 

5.  Can  combustion  take  place  without  the  evolution  of  light? 

6.  Is  the  evolution  of  light  always  produced  by  combustion  ? 

7.  Suggest  a  reason  why  wood  in  the  form  of  shavings  burns 
more  rapidly  than  the  same  wood  in  the  form  of  a  log. 

8.  Why  do  substances  burn  more  rapidly  in  pure  oxygen  than 
in  air? 

9.  Inquire  at  a  garage  whether  oxygen  is  used  in  connection 
with  the  repair  of  motor  cars  and,  if  so,  for  what  purpose. 

10.  Suggest  a  reason  why  some  metals  tarnish  on  exposure  to 
air,  while  others  remain  bright. 

11.  Consult  the  dictionary  for  the  derivation  and  significance  of 
the  word  phlogiston. 

12.  Calculate  the  weight  in  grams  of  50  1.  of  oxygen  measured 
at,  0°  and  under  a  pressure  of  1  atmosphere. 


OXYGEN  37 

13.  Calculate  the  weight  in  grams  of  each  of  the  following  com- 
pounds necessary  for  the  preparation  of  50  1.  of  oxygen  :  (1)  mercuric 
oxide ;  (2)  water ;  (3)  potassium  chlorate. 

14.  Calculate  the  volume  of  oxygen  that  would  be  evolved  on 
heating  10  g.  of  potassium  chlorate  ;  10  g.  of  mercuric  oxide. 

15.  Assuming  the  cost  of  potassium  chlorate  and  mercuric  oxide 
to  be  respectively  $0.50  and  $1.50  per  kilogram,  calculate  the  cost  of 
materials  necessary  for  the  preparation  of  50  1.  of  oxygen  from  each 
of  the  above  compounds. 

16.  100  g.  of  potassium  chlorate  and  25  g.  of  manganese  dioxide 
were  heated  in  the  preparation  of  oxygen.    Give  the  weight  of  each 
of  the  products  left  in  the  flask. 


CHAPTER  IV 
HYDROGEN 

Historical.  The  element  hydrogen  was  first  clearly 
recognized  as  a  distinct  substance  by  the  English  inves- 
tigator Cavendish,  who  in  1766  obtained  it  in  a  pure  state 
and  showed  it  to  be  different  from  the  other  inflammable 
airs  or  gases  which  had  long  been  known.  Because  it  had 
been  found  to  be  a  constituent  of  water,  Lavoisier  gave  it 
the  name  hydrogen,  which  means  "  water  former." 

Occurrence.  In  the  free  state  hydrogen  is  found  in  the 
atmosphere,  but  only  in  traces  (about  1  volume  in  from 
15,000  to  20,000  volumes  of  air).  In  the  combined  state  it 
is  widely  distributed,  being  a  constituent  of  water  as  well 
as  of  all  living  organisms  and  the  products  derived  from 
them,  such  as  starch  and  sugar.  About  10  per  cent  of 
the  human  body  is  hydrogen.  Combined  with  carbon,  it 
forms  the  substances  which  constitute  petroleum  and 
natural  gas.  It  is  an  interesting  fact  that  while  hydro- 
gen in  the  free  state  occurs  only  in  traces  on  the  earth, 
it  occurs  in  enormous  quantities  in  the  gaseous  matter 
surrounding  the  sun  and  certain  other  stars. 

Preparation  from  water.  Hydrogen  can  be  prepared 
from  water  by  several  methods,  the  most  important  of 
which  are  the  following : 

1.  By  the  electric  current.  As  has  been  indicated  in 
Chapter  II,  water  is  easily  separated  into  its  constituents, 
hydrogen  and  oxygen,  by  passing  an  electric  current 
through  it  under  certain  conditions. 


HYDROGEN 


39 


2.  By  the  action  of  certain  metals.  When  brought  into 
contact  with  certain  metals  under  appropriate  conditions, 
water  gives  up  a  part  or  the  whole  of  the  hydrogen,  the 
place  of  the  hydrogen  being  taken 
by  the  metal.  In  the  case  of  a  few 
of  the  metals  this  change  occurs  at 
ordinary  temperatures.  Thus,  if  a 
bit  of  sodium  is  thrown  on  water, 
an  action  is  seen  to  take  place  at 
once,  sufficient  heat  being  gener- 
ated to  melt  the  sodium,  which 
runs  about  on  the  surface  of  the 
water.  The  change  which  takes 
place  consists  in  the  displacement 
of  one  half  of  the  hydrogen  of 
the  water  by  the  sodium,  and  may 
be  represented  as  follows: 

sodium  +  water >-  hydrogen  +  sodium  hydroxide 

["hydrogen")  fsodium     ~| 


FIG.  14.    The  preparation 

of  hydrogen  by  the  action 

of  sodium  on  water 


[oxygen     J 


L  hydrogen J 
oxygen     J 


The  sodium  hydroxide  formed  is  a  white  solid  which 
remains  dissolved  in  the  excess  of  undecomposed  water 
and  may  be  obtained  by  evaporating  the  solution  to  dry- 
ness.  The  hydrogen  is  evolved  as  a  gas  and  may  be 
collected  by  suitable  means. 

A  simple  form  of  apparatus  used  in  preparing  hydrogen  by 
the  action  of  sodium  on  water  is  represented  in  Fig.  14.  Since 
the  sodium  is  lighter  than  water,  it  is  kept  under  the  water  by 
pushing  a  pellet  of  the  metal  into  the  end  of  a  short  piece  of 
lead  tubing,  the  other  end  of  which  has  been  sealed.  The  tube 
containing  the  sodium  is  then  dropped  into  a  trough  of  water. 
Hydrogen  is  at  once  evolved  and  is  collected  by  bringing  over 
it  a  bottle  or  cylinder  filled  with  water,  as  shown  in  the  figure. 


40     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Other  metals,  such  as  magnesium  and  iron,  decompose 
water  rapidly  but  only  at  higher  temperatures.  When 
steam  is  passed  over  hot  iron,  for  example,  the  iron  com- 
bines with  the  oxygen  of  the  steam,  setting  free  all  of 
the  hydrogen.  Experiments  show  that  the  change  may 
be  represented  as  follows: 

iron  +  water >-  hydrogen  +  iron  oxide 

[hydrogen]  [iron      "I 

oxygen     J  [oxygenj 

The  iron  oxide  formed  is  a  reddish-black  compound  iden- 
tical with  that  obtained  by  the  combustion  of  iron  in  oxygen. 

Preparation  of  hydrogen  from  iron  and  steam.  The  apparatus 
used  in  the  preparation  of  hydrogen  from  iron  and  steam  is 
shown  in  Fig.  15.  A  porcelain  or  iron  tube  A,  about  50  cm.  in 
length  and  2  cm.  or  3  cm.  in  diameter,  is  partly  filled  with  fine 
iron  wire  or  tacks  and  connected  as  shown  in  the  figure.  The 
tube  is  heated,  slowly  at  first,  until  the  iron  is  red-hot.  Steam 
is  then  conducted  through  the  tube  by  boiling  the  water  in  the 
flask  B.  The  hot  iron  combines  with  the  oxygen  in  the  steam, 
setting  free  the  hydrogen,  which  is  collected  over  water  in  C. 

Preparation  from  acids.  A  more  convenient  method  for 
preparing  hydrogen  in  the  laboratory  consists  in  liberating 
it  from  acids  by  the  action  of  metals.  For  this  purpose 
any  of  the  metals  which  liberate  hydrogen  from  water, 
but  only  these,  may  be  employed.  Usually  zinc  or  iron 
is  used.  The  acids  commonly  employed  are  either  hydro- 
chloric acid  or  sulfuric  acid.  The  former  is  an  aqueous 
solution  of  a  gaseous  compound  (known  as  hydrogen 
chloride)  which  contains  2.76  per  cent  of  hydrogen  and 
97.24  per  cent  of  chlorine,  while  the  latter  is  an  aqueous 
solution  of  an  oily  liquid  (known  as  hydrogen  sulfate) 
which  consists  of  2.06  per  cent  of  hydrogen,  32.69  per 
cent  of  sulfur,  and  65.25  per  cent  of  oxygen.  To  liberate 


HYDROGEN 


41 


hydrogen  it  is  only  necessary  to  bring  the  acid,  properly 
diluted  with  water,  into  contact  with  the  metal.  The 
metal  gradually  passes  into  solution,  while  the  hydrogen  of 
the  acid  is  in  turn  set  free.  The  liberation  of  the  hydro- 
gen is  indicated  by  the  effervescence  of  the  liquid.  When 


FIG.  15.   The  preparation  of  hydrogen  by  the  action  of  iron  on  steam 

zinc  and  sulfuric  acid  are  used  in  the  preparation,  the 
reaction  may  be  represented  in  a  general  way  as  follows: 

zinc  +  sulfuric  acid >•  zinc  sulfate  +  hydrogen 

[hydrogen!  fzinc      ~| 

sulfur  sulfur 

oxygen    J  LoxygenJ 

It  will  be  noted  that  the  zinc  simply  takes  the  place  of 
the  hydrogen  in  the  acid.  The  resulting  compound  of 
zinc,  sulfur,  and  oxygen,  known  as  zinc  sulfate,  is  a  white 
solid  which  remains  dissolved  in  the  water  present. 

When  iron  and  hydrochloric  acid  are  used  in  the  prepara- 
tion of  hydrogen,  the  reaction  may  be  represented  as  follows : 


iron  +  hydrochloric  acid 


rhydrogen-l 
[chlorine    J 


iron  chloride  +  hydrogen 

[iron        1 
(.chlorine  J 


42     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


FIG.  16.    The  preparation  of  hydrogen  by  the 
action  of  metals  on  acids 


Laboratory  apparatus.  A  convenient  form  of  apparatus  for 
preparing  hydrogen  by  the  action  of  metals  upon  acids  is 
shown  in  Fig.  16.  The  metal  is  placed  in 
flask  A,  which  is  fitted  with  a  stopper  and 
connected  with  tubes,  as  shown  in  the  fig- 
ure. The  acid,  properly  diluted  with  water, 
is  added  a  little  at  a  time  through  the  funnel 
tube  B.  The  liber- 
ated hydrogen  es- 
capes through  C 
and  is  collected 
in  receivers,  as 
shown  in  the  fig- 
ure. The  hydrogen 
which  first  escapes 
through  the  exit 
tube  is  mixed  with  the  air  originally  present  in  flask  A.  Such 
a  mixture  of  hydrogen  and  air  is  violently  explosive  when 
brought  in  contact  with  a  flame.  There- 
fore one  must  keep  all  flames  away  from 
the  apparatus.  Moreover,  one  should  not 
collect  the  hydrogen  until  an  amount  of  it 
has  been  generated  sufficient  to  displace 
all  the  air  previously  contained  in  the  flask. 
A  more  convenient  form  of  apparatus 
to  use  is  that  shown  in  Fig.  17.  It  is 
known  as  a  Kipp  generator  and  it  has  the 
advantage  of  being  automatic  in  its  ac- 
tion. The  metal  is  placed  in  A,  and  the 
acid  poured  into  B.  When  the  stopcock 
D  is  opened,  the  acid  runs  down  into  C 
and  up  into  A,  where  it  comes  in  contact  FIG  17  A  R. 
with  the  metal.  The  hydrogen  generated  erator  for  preparing 
escapes  through  D.  If  now  the  stopcock  hydrogen 

is  closed,  the  hydrogen,  being  unable  to 
escape  through  the  tube,  pushes  the  aoid  away  from  the  metal 
in  A  down  into  C  and  up  into  .B,  so  that  the  action  ceases. 


HYDKOGEN 


43 


Commercial  method  for  preparing  hydrogen.  A  number 
of  processes  are  being  used  for  preparing  hydrogen  on  a 
large  scale,  and  new  methods  are  being  investigated  in 
the  hope  of  finding  still  cheaper  ones.  At  present,  in 
the  United  States,  it  is  prepared  from  water,  either  by 
passing  steam  over  hot  iron  (p.  40)  or  by  decomposing 
the  water  by  an  electric  current  (p.  16).  Another  method, 
which  is  used  in  Europe,  consists  in  distilling  petroleum 
under  definite  conditions.  Petroleum  is  composed  largely 
of  compounds  of  hydrogen  and  carbon,  and  when  it  is 
heated  under  proper  conditions  hydro- 
gen is  set  free. 

Properties  of  hydrogen.  Hydrogen, 
like  oxygen,  is  a  colorless,  odorless, 
and  tasteless  gas.  One  liter  of  it 
weighs  0.08987  g.  It  is  the  lightest 
of  all  known  substances,  being  14.385 
times  lighter  than  air;  it  may  there- 
fore be  transferred  from  one  vessel 

,,        ,  .,  j  drogen    upward    from 

to  another  by  pouring  it  upward,  as     0  ^Vessel  into  ^other 

shown  in  Fig.  18.  The  hydrogen  in  the 
cylinder  A  rises  to  the  top  of  the  cylinder  B  and  forces 
the  air  out.  The  solubility  of  hydrogen  in  water  is  very 
small,  being  only  about  one  half  as  great  as  that  of  oxygen. 
Dewar  was  the  first  to  obtain  hydrogen  in  the  liquid 
state.  He  cooled  the  gas  to  a  temperature  of  —  205°  by 
means  of  liquid  air,  and  at  the  same  time  subjected  it 
to  a  pressure  of  180  atmospheres.  It  was  obtained  as  a 
colorless,  transparent  liquid,  boiling  at  —  252.7°.  This  is 
the  lightest  liquid  known,  having  a  density  of  but  0.07  at 
its  boiling  point.  When  liquid  hydrogen  is  evaporated 
under  very  small  pressure,  solid  hydrogen  is  obtained  as 
a  snowlike  body  melting  at  about  —  259°. 


FIG.  18.     Pouring   hy- 


44     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

A  number  of  the  metals  have  the  property  of  absorbing, 
or  occluding,  hydrogen.  Most  of  the  metals  absorb  only 
a  small  volume  of  the  gas,  but  a  few,  such  as  gold,  plati- 
num, and  palladium,  absorb  large  volumes.  One  volume 
of  finely  powdered  palladium  absorbs  over  800  volumes 
of  the  gas  at  ordinary  temperatures. 

Chemical  conduct.  At  ordinary  temperatures  hydrogen 
is  not  an  active  element.  Under  suitable  conditions,  how- 
ever, it  combines  with  many  of  the  elements,  forming 
compounds  known  as  hydrides.  Thus,  hydrogen  and 
chlorine,  when  mixed  together,  will  combine  with  ex- 
plosive violence  if  heated  or  if  exposed  to  the  sunlight. 
The  hydride  formed  in  either  case  is  called  hydrogen 
chloride.  Under  suitable  conditions  hydrogen  combines 
with  nitrogen  to  form  ammonia,  and  with  sulfur  to  form 
the  foul-smelling  gas  hydrogen  sulfide.  At  ordinary  tem- 
peratures hydrogen  and  oxygen  may  be  mixed  without 
apparent  action.  If  the  mixture  is  heated  to  about  800°, 
or  if  a  flame  is  brought  in  contact  with  it,  a  violent 
explosjon  takes  place.  Nevertheless,  under  proper  con- 
ditions hydrogen  may  be  made  to  burn  quietly  in  either 
oxygen  or  air.  The  resulting  hydrogen  flame  is  almost 
colorless  and  is  very  hot.  The  combustion  of  the  hydrogen 
is  due  to  its  union  with  oxygen,  and  the  product  of  the 
combustion  is  water.  Experiments  show  that  the  ratio  in 
which  the  two  gases  combine  is  1  part  of  hydrogen  to 
7.94  of  oxygen  by  weight.  The  heat  liberated  in  the 
reaction  amounts  to  34,226  cal.  for  each  gram  of  hydrogen 
entering  into  combination. 

Directions  for  burning  hydrogen.  The  combustion  of  hydrogen 
in  air  may  be  carried  out  safely  as  follows :  The  hydrogen 
is  generated  in  the  bottle  A  (Fig.  19),  is  dried  by  conducting 
it  through  the  tube  B  filled  with  some  substance  (usually 


HYDROGEN  45 

calcium  chloride)  which  has  a  great  attraction  for  moisture,  and 
escapes  through  the  tube  C,  the  end  of  which  is  drawn  out  to 
a  jet.  When  all  the  air  has  been  expelled  from  the  apparatus 
the  hydrogen  may  be  ignited.  It  then  burns 'quietly,  since  only 
the  small  amount  of  it  which  escapes  from  the  jet  can  come 
in  contact  with  the  oxygen  of  the  air  at  any  one  time.  By 
holding  a  cold  dry  bell  jar  or  bottle  D  over  the  flame  in  the 


FIG.  19.   Burning  hydrogen  and  collecting  the  product  of  its  combustion 

manner  shown  in  the  figure,  the  steam  formed  by  the  combus- 
tion of  the  hydrogen  is  condensed,  and  water  collects  in  drops 
on  the  sides  of  the  jar. 

Temperature  at  which  hydrogen  combines  with  oxygen.  The 
union  of  hydrogen  and  oxygen  probably -takes  place  at  ordinary 
temperatures,  but  the  speed  of  the  reaction  is  so  slow  that  no 
combination  can  be  detected  even  after  long  intervals  of  time. 
As  the  temperature  is  raised  the  speed  increases.  Thus,  Meyer 
and  Raum  found  that  the  two  gases,  when  mixed  in  the  pro- 
portion of  two  volumes  of  hydrogen  to  one  volume  of  oxygen 
and  heated  to  100°  for  218  days,  showed  no  appreciable  com- 
bination. When  heated  to  300°  for  65  days  it  was  found  that, 
in  different  trials,  from  0.4  per  cent  to  9.5  per  cent  of  the  mix- 
ture had  combined.  At  500°  the  change  is  still  more  rapid, 
but  requires  several  hours  for  completion.  At  a  temperature 
roughly  approximating  800°  the  union  of  the  two  takes  place 


46     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


FIG.  20.    Exploding  a 

mixture    of    hydrogen 

and  oxygen 


with  explosive  violence.    The  temperature  at  which  this  in- 
stantaneous combination  takes  place  is  modified  by  very  slight 
changes  in  the  conditions,  due  to  the  catalytic  effect  of  foreign 
substances,    such  as  moisture  and  the 
ji  materials  of  which  the  tube  containing 

JraL^  the  gases  is  made.   Certain  catalyzers, 

such  as  finely  divided  platinum,  bring 
about  practically  instantaneous  combi- 
nation at  ordinary  temperatures. 

A  mixture  of  hydrogen  and  oxygen  is 
explosive.  That  a  mixture  of  hydrogen 
and  air  is  explosive  may  be  shown 
safely  as  follows  :  A  cork  through  which 
passes  a  short  glass  tube  about  1  cm. 
in  diameter  is  fitted  air-tight  into  the 
tubule  of  a  bell  jar  (Fig.  20)  of  2  1.  or  3  1. 

capacity.  (A  thick  glass  bottle  with  the  bottom  removed  may 
be  used.)  The  tube  is  closed  with  a  small  rubber 
stopper  and  the  bell  jar  filled  with  hydrogen, 
the  gas  being  collected  over  water.  When  en- 
tirely filled  with  the  gas  the  jar  is  removed  from 
the  water  and  supported  by  blocks  of  wood  in 
order  to  leave  the  bottom  open,  as  shown  in 
Fig.  20.  The  stopper  is  now  removed  from  the 
tube  in  the  cork.  The  hydrogen,  on  account  of 
its  lightness,  escapes  from  the  tube  and  is  at 
once  lighted.  As  the  hydrogen  escapes,  the 
air  flows  in  at  the  bottom  of  the  jar  and  mixes 
with  the  remaining  portion  of  the  hydrogen, 
so  that  a  mixture  of  the  two  soon  forms,  and  a 
loud  explosion  results.  The  explosion  is  not 
dangerous,  since  the  bottom  of  the  jar  is  open, 
thus  leaving  room  for  the  expansion  of  the 
hot  gas. 

Since  air  is  only  one  fifth  oxygen,  the  remainder  being  inert 
gases,  it  may  readily  be  inferred  that  a  mixture  of  hydrogen 
with  pure  oxygen  would  be  far  more  explosive  than  a  mixture 


FIG.  21.  Flame 
of  a  candle  ex- 
tinguished    by 
hydrogen 


HYDROGEN 


47 


of  hydrogen  with  air.    Such  mixtures  should  not  be  made 
except  in  small  quantities  and  by  experienced  workers. 

Hydrogen  does  not  support  combustion.  While  hydrogen 
is  readily  combustible,  it  is  not  a  good  supporter  of  com- 
bustion ;  in  other  words,  most  substances  will  not  burn  in 
it.  This  may  be  shown  by  bringing  a  lighted  candle  sup- 
ported by  a  stiff  wire  into  a  bottle  or  cylinder  of  the  pure 
gas,  as  shown  in  Fig.  21.  The  hydrogen  is  ignited  by  the 
flame  of  the  candle  and  burns  at  the  mouth  of  the  cylinder, 
where  it  comes  in 
contact  with  the  •* 
oxygen  in  the  air. 
When  the  candle 
is  thrust  up  into 
the  gas,  its  flame 
is  extinguished. 
If  slowly  with- 
drawn, the  candle 
is  relighted  as  it 
passes  through 
the  layer  of  burn- 
ing hydrogen. 

Reduction.  On  account  of  its  tendency  to  combine  with 
oxygen,  hydrogen  has  the  power  of  abstracting  it  from  many 
of  its  compounds.  Thus,  if  a  stream  of  hydrogen  generated 
in  A  (Fig.  22)  and  dried  by  passing  through  the  tube  B 
(filled  with  calcium  chloride)  is  conducted  through  the  tube 
(7,  which  contains  some  copper  oxide  heated  to  a  moderate 
temperature,  the  hydrogen  abstracts  the  oxygen  from  the 
copper  oxide.  The  change  may  be  represented  as  follows : 

copper  oxide  +  hydrogen >•  water  -f  copper 


FIG.  22.    The  reduction  of  hot  copper  oxide  by 
a  stream  of  hydrogen 


[copper  1 
LoxygenJ 


[hydrogen! 
Loxygen     J 


48     AN  ELEMENTABY  STUDY  OF  CHEMISTRY 


The  oxides  of  most  of  the  metals  act  in  a  similar  way. 
Thus,  when  hydrogen  is  passed  over  hot  iron  oxide,  water 
and  iron  are  formed: 

iron  oxide  +  hydrogen >•  water  +  iron 

[iron      1  rhydrogenl 

oxygenj  Loxygen     J 

In  these  reactions  the  oxide  of  the  metal  is  said  to 
undergo  reduction.  Reduction  may  be  denned,  therefore, 

as  the  process  of  removing 
oxygen  from  a  compound. 
An  element,  such  as  hydro- 
gen, which  has  a  strong  affin- 
ity for  oxygen  and  which 
may  be  used  for  removing 
oxygen  from  a  compound 
is  termed  a  reducing  agent. 
Relation  of  oxidation  and 
reduction.  It  is  evident 
from  the  statements  con- 
cerning oxidation  and  reduc- 
tion that  the  two  processes 
are  just  the  opposite  of 
each  other.  The  one  process 
consists  in  the  addition  of 
oxygen  to  an  element  or 
compound,  while  the  other 
consists  in  the  removal  of 
oxygen  from  a  compound. 


FIG.  23.    Robert  Hare  (1781-1858) 

An  early  American  chemist ;  the  in- 
ventor of  many  ingenious  laboratory 
appliances,     including    the     oxyhy- 
drogen  blowpipe 


Moreover  it  usually  happens 
that  when  one  substance  is  oxidized,  some  other  substance 
is  reduced.  Thus,  when  hydrogen  is  passed  over  hot  copper 
oxide  (Fig.  22),  the  hydrogen  is  oxidized,  while  the  copper 
oxide  is  reduced.  It  will  be  pointed  out  later  that  the 


HYDROGEN  49 

terms  oxidation  and  reduction  are  sometimes  used  with  a 
broader  meaning  than  those  just  given. 

The  oxyhydrogen  blowpipe.  This  is  a  form  of  apparatus  used 
for  burning  hydrogen  in  pure  oxygen.  It  was  devised  and 
first  used  by  an  American  scientist  Robert  Hare  (Fig.  23)  in 
the  year  1801,  in  his  laboratory  in  Philadelphia.  As  has  been 
previously  stated,  the  flame  produced,  by  the  combustion  of 
hydrogen  in  the  air  is  very  hot.  It  is  evident  that  if  pure 
oxygen  is  substituted  for  air,  the  temperature  reached  will  be 
much  higher,  since  there  are  no  inert  gases  to  absorb  the  heat. 
The  oxyhydrogen  blowpipe, 
used  to  effect  this  combina- 
tion, consists  of  a  small  tube 
placed  within  a  larger  one, 
as  shown  in  Fig.  24. 

The  hydrogen,  stored  un- 
der pressure,  usually  in  steel 
cylinders,  is  first  passed 
through  the  outer  tube  // 
(Fig.  24)  and  ignited  at  the 
open  end  of  the  tube  A.  The  FlG.  24.  The  oxyhydrogen  blowpipe 
oxygen  from  a  similar  cylin- 
der is  then  conducted  through  the  inner  tube  0,  and  mixes 
with  the  hydrogen  at  the  end  of  the  tube.  In  order  to  produce 
the  maximum  heat,  the  hydrogen  and  oxygen  must  be  admitted 
to  the  blowpipe  in  the  exact  proportion  in  which  they  com- 
bine ;  namely,  2  volumes  of  hydrogen  to  1  of  oxygen,  or,  by 
weight,  1  part  of  hydrogen  to  7.94  parts  of  oxygen.  The  inten- 
sity of  the  heat  may  be  shown  by  bringing  into  the  flame 
pieces  of  metal  such  as  iron  wire  or  zinc.  These  burn  with 
great  brilliancy.  Even  platinum,  which  has  a  melting  point 
of  1755°,  may  be  melted  by  the  heat  of  the  flame. 

While  the  oxyhydrogen  flame  is  intensely  hot,  it  is  almost 
nonluminous.  If  the  flame  is  directed  against  some  infusible 
substance  like  ordinary  lime  (calcium  oxide),  the  heat  is  so 
intense  that  the  lime  becomes  incandescent  and  glows  with 


50     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

a  brilliant  light.  This  is  sometimes  used  as  a  source  of  light, 
under  the  name  of  Drummond  light  or  limelight. 

The  blast  lamp.  A  similar  form  of  apparatus  is  used  in  the 
laboratory  as  a  source  of  heat  under  the  name  blast  lamp 
(Fig.  25).  This  differs  from  the  oxy hydrogen  blowpipe  only 
in  the  size  of  the  tubes.  In  place  of  the  hydrogen  and  oxygen 
the  more  accessible  coal  gas  (or  natural  gas)  and  air  are  used. 
Coal  gas  and  natural  gas  are  composed  largely  of  a  mixture  of 
gaseous  compounds  of  carbon  and  hydrogen  (Chap.  XXV). 
While  the  temperature  of  the  flame  is  not  so  high  as  that  of 
the  oxyhydrogen  blowpipe,  it 
nevertheless  suffices  for  most 
chemical  operations. 

Exothermic    and    endothermic 
reactions.    We    have   seen    that 
in     certain     reactions     a     large 
amount    of    heat    is    given    off. 
Examples  of  such  reactions  are 
the  combustion   of  hydrogen  in 
oxygen  and  the   combustion  of 
FIG.  25.   The  ordinary  labora-     CQal  in   air>     Au   guch   reactiOns 
tory  blast  lamp  . 

are    said    to    be    exothermic.     In 

other  reactions  heat  is  continuously  absorbed  and  so  must 
be  applied  in  order  that  the  reaction  may  continue.  Such 
reactions  are  termed  endothermic.  The  decomposition  of 
sugar  is  an  example  of  an  endothermic  reaction. 

Uses  of  hydrogen.  Hydrogen  is  used  as  a  material  for 
the  inflation  of  dirigible  balloons  such  as  are  used  in  war. 
It  has  been  used  in  the  oxyhydrogen  blowpipe  as  a  source 
of  light  and  heat,  but  this  use  has  largely  given  way  to 
more  economical  methods  of  producing  high  temperatures. 
(See  the  oxy  acetylene  blowpipe  and  the  electric  furnace.) 
It  is  also  used  as  an  agent  for  purifying  certain  oils,  and 
for  converting  these  oils  into  solid  fats. 


HYDROGEN  51 

EXERCISES 

1.  Report  the  important  events  in  the  life  of   Cavendish  (con- 
sult encyclopedia). 

2.  Will  a  definite  weight  of  iron  decompose  an  unlimited  weight 
of  steam  ? 

3.  Calculate  the  relative  weight  of  hydrogen  and  oxygen  (see 
weights  of  1 1.  of  different  gases,  in  Appendix). 

4.  Why  is  oxygen  passed  through  the  inner  tube  of  the  oxyhy- 
drogen  blowpipe  rather  than  the  outer  ? 

5.  What  is  the  source  of  heat  in  the  limelight? 

6.  In  the   experiment   illustrated  by  Fig.  21,  will   the    flame 
remain  at  the  mouth  of  the  cylinder  ? 

7.  Distinguish    clearly  between  the  following  terms:  oxidation, 
reduction,  combustion,  and  kindling  temperature. 

8.  Is  oxidation  always  accompanied  by  reduction  ? 

9.  In  the  experiment  illustrated  by  Fig.  19,  why  dry  the  hydro- 
gen before  burning  it  ? 

10.  Suggest  a  way  of  determining  the  weight  of  the  water  formed 
in  the  reaction  in  the  experiment  illustrated  by  Fig.  22. 

11.  If  hydrogen  and  oxygen  unite  in  the  ratio  of  1  to  7.94  by 
weight,  in  what  ratio  do  they  unite  by  volume  ? 

12.  (a)  How  many  calories  of  heat  are  evolved  in  the  combustion 
of  100  1.  of  hydrogen  to  form  water?  (fe)  How  many  grams  of  water 
are  formed? 

13.  How  many  grams  of  hydrogen  can  be  obtained  from  100  g. 
of  hydrogen  sulf ate  ?   What  volume  would  this  amount  of  hydrogen 
occupy  ? 

14.  A  gas  tank  holds  250  1.  of  hydrogen,    (a)  What  is  the  weight 
of  this  volume  of  hydrogen?    (ft)  What  weight  of  water  would  have 
to  be  decomposed  by  electricity  in  order  to  prepare  this  volume  of 
hydrogen?    (c)  What  volume  of  oxygen  would  be  liberated  in  the 
process? 

15.  10  g.  of  water  is  boiled  and  the  steam  passed  over  heated 
iron  (Fig.   15).    (a)  What  weight  of  hydrogen  will  be  liberated  ? 
(6)  What  volume  will  the  hydrogen  occupy  ?    (c)  What  change  will 
take  place  in  the  weight  of  the  iron  ? 


CHAPTER  V 


THE  GAS  LAWS;    THE  KINETIC  THEORY 

Introduction.  It  will  be  remembered  that  in  describing 
the  properties  of  oxygen  and  hydrogen  the  weight  of  a 
liter  of  each  gas  was  given.  A  moment's  reflection  will 

make  it  clear  that  these 
weights  must  be  correct 
only  under  certain  condi- 
tions, for  it  is  a  familiar 
fact  that  the  volume  of  a 
given  weight  of  a  gas  varies 
both  with  changes  in  pres- 
sure and  with  changes  in 
temperature. 

Variation  of  volume  with 
pressure :  law  of  Boyle. 
That  the  volume  occupied 
by  a  given  weight  of  gas 
can  be  altered  by  changing 
the  pressure  is  familiar  to 
everyone  who  has  pumped 
air  into  a  bicycle  or  auto- 
mobile tire.  As  early  as 
1660  Robert  Boyle,  an  Irish  investigator  (Fig.  26),  showed 
that  the  following  statement  correctly  expresses  the  rela- 
tion between  volume  and  pressure.  If  the  temperature  re-, 
mains  constant,  the  volume  occupied  by  a  given  weight  of  a 
gas  varies  inversely  as  the  pressure  to  which  it  is  subjected. 


FIG.  26.   Robert  Boyle  (1627-1691) 

One  of  the  most  accurate  of  the  early 
experimenters  in  chemistry  and  physics 


THE  GAS  LAWS;  THE  KINETIC  THEORY      53 

This  generalization  is  known  as  Boyle's  law.  Thus,  if  a 
given  weight  of  a  gas  occupies  a  volume  of  1000  cc.  when 
subjected  to  a  certain  pressure,  it  will  occupy  a  volume 
of  500  cc.  if  the  pressure  is  doubled,  or  of  250  cc.  if  the 
pressure  is  made  four  times  as  great,  or  of  2000  cc.  if  the 
pressure  is  diminished  one  half.  This  means  that  for  a 
given  weight  of  gas  the  product  of  the  pressure  by  the  vol- 
ume will  remain  constant,  no  matter  how  either  one  may 
be  altered.  Designating  the  pressure  and  volume  under 
one  set  of  conditions  by  P  and  F,  and  under  a  different 
set  by  Pl  and  Fx,  Boyle's  law  may  be  stated  thus: 


Standard  pressure.  For  practical  purposes  we  must 
choose  some  standard  pressure  to  which  all  gas  volumes 
may  be  referred.  This  is  most  conveniently  chosen  as  the 
average  pressure  of  the  atmosphere  at  the  sea  level.  This 
is  equal  to  1033.3  g.  per  square  centimeter.  In  place  of 
expressing  the  pressure  in  this  way  it  is  much  more  con- 
venient to  express  it  in  terms  of  the  height  of  the  column 
of  mercury  which  the  pressure  of  the  atmosphere  will 
sustain.  Expressed  in  this  way  the  standard  pressure  is 
equal  to  that  exerted  by  a  column  of  mercury  760  mm. 
in  height  —  this  being  the  average  height  of  the  barometer 
at  the  sea  level. 

Illustration  of  the  law  of  Boyle.  A  typical  example  will  make 
the  meaning  of  this  law  clear.  Suppose  that  a  gas,  measured 
under  a  pressure  of  720  mm.,  has  a  volume  of  620  cc.  What 
volume  will  this  gas  occupy  under  standard  pressure  (760  mm.), 
the  temperature  remaining  constant  ? 

According  to  Boyle's  law,  PF=P1F1.  Substituting  the 
values  given  in  the  problem,  we  have  760  x  F  =  720  x  620 ; 
or  V  =  587.4  cc. 


54     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Centigrade 


100° 


Variation  of  volume  with  temperature.  If  the  pressure 
is  held  constant,  all  gases  expand  when  the  temperature 
is  raised  and  contract  when  it  is  lowered,  and  it  is  a 
remarkable  fact  that  the  volumes  of  all  gases  change  to 
the  same  extent  for  a  given  varia- 
tion in  the  temperature.  Let  us 
suppose  that  the  volume  of  a  gas 
has  been  measured  at  zero  on 
the  centigrade  scale.  Experiment 
has  shown  that  a  rise  of  one  degree 
causes  an  expansion  of  ^L  of  this 
volume ;  a  rise  of  five  degrees,  an  ex- 
pansion of  ^Jy.  If  we  take  273  cc. 
of  this  gas  at  zero,  the  volume  at  1° 
above  will  be  274  cc.;  at  1°  below 
it  will  be  272  cc. ;  and  at  5°  below 
it  will  be  268  cc.  At  the  same  rate 
of  contraction  the  volume  will  be 
1  cc.  at  -  272°,  and  at  -  273°  it 
will  be  zero.  Of  course  this  can- 
not really  happen,  and  experiment 
shows  that  before  this  temperature 
is  reached,  all  gases  have  changed 
into  liquids  or  solids.  Helium,  the 
most  difficult  gas  to  liquefy,  passes 
into  a  liquid  at  -  268.7°. 

The  absolute  scale  of  temperature.  If  we  were  to  con- 
struct a  thermometer  having  divisions  of  the  same  size 
as  those  on  the  centigrade  scale,  but  with  the  zero  point 
at  —  273°  on  the  latter  scale,  then  the  point  at  which 
water  freezes  (0°  centigrade)  would  be  273°.  At  272° 
on  this  scale  the  273  cc.  of  gas  mentioned  in  the  last  para- 
graph would  measure  272  cc. ;  at  271°  it  would  measure 


-£52.7' 
-273° 


Freezing  Point 
of  Water   '" 


...Boiling  Point__ 
o/'  Ozone 


Boiling  Point 
"•of  Hydrogen  — 
--Absolute  Zero— 


FIG.  27.    Comparison  of  the 

centigrade  with  the  absolute 

scale  of  temperature 


THE  GAS  LAWS;  THE  KINETIC  THEORY      55 


271  cc. ;  at  1°,  1  cc.  On  such  a  scale  the  volume  of  a  gas 
would  be  proportional  to  the  temperature  at  every  point. 
This  scale  is  known  as  the  scale  of  absolute  temperature,  the 
point  —  273°  centigrade  being  the  absolute-zero  point.  Evi- 
dently the  absolute  temperature  may  be  obtained  by  adding 
273°  to  the  centigrade  read- 
ing. Thus  30°  centigrade 
(C.)  equals  30°  +  273° 
(or  303°)  absolute  (A.). 
Fig.  27  gives  a  comparison 
of  the  centigrade  and  abso- 
lute scales  at  a  number  of 
temperatures. 

The  law  of  Gay-Lussac 
(or  of  Charles).  A  general 
statement  can  now  be  made 
in  regard  to  the  effect  of 
temperature  on  the  volume 
of  a  gas :  If  the  pressure 
remains  constant,  the  vol- 
umes occupied  by  a  given 
weight  of  a  gas  at  differ- 
ent temperatures  are  directly 
proportional  to  the  absolute 
temperatures  of  the  gas. 

If  V  and  V1  are  the  volumes  at  the  temperatures  T  and 
TI  respectively,  then  we  have  the  following  proportion : 

Z_Z 
V~  2\ 

The  above  generalization  is  called  the  law  of  Gay-Lussac 
(Fig.  28)  or  of  Charles,  since  it  was  formulated  inde- 
pendently by  these  two  Frenchmen  in  1801. 


FIG.  28.    Joseph  Louis  Gay-Lussac 
(1778-1850) 

A  distinguished  French  chemist  who 
contributed  much  to  our  knowledge 
of  gases  and  their  combining  ratios 


56     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Illustration  of  the  law  of  Gay-Lussac.  The  following  example 
will  make  the  meaning  of  the  law  clear.  The  volume  of  a  cer- 
tain gas  measured  at  a  temperature  of  70°  is  650  cc.  What 
will  be  its  volume  at  10°  ? 

First  reduce  the  centigrade  readings  to  absolute  : 

70°  C.  =  70  +  273  =  343°  A.  ;  10°  C.  =  10  +  273  =  283°  A. 

Now  substituting  the  appropriate  values  in  the  above  equation, 
we  have  OCQ 


Variations  in  volume  due  to  changes  both  in  pressure 
and  temperature.  In  case  both  pressure  and  temperature 
change,  then  the  correction  may  be  made  for  each  in 
succession,  as  illustrated  in  the  following  example  : 

A  certain  weight  of  gas  measured  500  cc.  at  a  temperature 
of  100°  when  subjected  to  a  pressure  of  760  mm.  Calculate 
the  volume  which  this  gas  will  occupy  at  a  temperature  of  50° 
and  a  pressure  of  740  mm. 

First  make  the  correction  for  pressure  : 


740  x  V  =  760  x  500  ;  or  V  =  513.5  cc. 
Next  make  the  correction  for  temperature  : 
V       T  V         323 


Standard  conditions.  Since  the  volume  of  a  gas  varies 
with  both  temperature  and  pressure,  it  is  essential  that 
we  select  both  a  standard  temperature  and  a  standard 
pressure  to  which  all  gas  volumes  may  be  referred.  We 
have  already  noted  that  the  standard  pressure  adopted  is 
that  exerted  by  a  column  of  mercury  760  mm.  in  height. 
As  a  standard  temperature,  the  temperature  of  melting  ice 


THE  GAS  LAWS;  THE  KINETIC  THEORY      57 


is  chosen.  This  is  0°  centigrade  or  273°  absolute.  When- 
ever the  volume  of  a  gas  is  given,  unless  otherwise  speci- 
fied, it  is  always  assumed  that  the  volume  given  is  that 
occupied  by  the  gas  under  standard  conditions. 

Standard  conditions  and  laboratory  conditions.  The  con- 
ditions of  temperature  and  pressure  which  prevail  in  the 
laboratory  are  never  the  standard  conditions.  Knowing 
the  volume  of  a  gas  under  laboratory  conditions,  however,  it 
is  a  simple  matter  to  calculate  the  vol- 
ume which  the  gas  will  occupy  under 
standard  conditions.  The  following 
problem  will  illustrate  the  method : 


A  gas  measured  300  cc.  under  a  pres- 
sure of  740  mm.  and  a  temperature  of 
25°  (or  298°  A.).  What  will  its  volume  be 
under  standard  conditions  (0°  and  760  mm. 
pressure)  ? 

First  find  the  change  in  volume  due  to 
change  in  pressure : 

300  x  740  =  760  x  V ;  or  V  =  292  cc. 

Next  make  the  correction  for  temper- 
ature : 

V         273 


Fi-o.29.  Measuringthe 
volume  of  a  gas  col- 
lected over  a  liquid 


Aqueous  vapor  pressure.  As  a  rule  gases  are  measured 
in  the  laboratory  by  collecting  them  over  water  in  a 
graduated  tube  as  shown  in  Fig.  29.  For  example,  if  we 
wish  to  measure  the  volume  of  oxygen  evolved  in  a  cer- 
tain reaction,  the  gas  is  conducted  into  the  tube  A,  which 
has  been  previously  filled  with  water  and  inverted  in  a 
cylinder  of  water  as  shown  in  the  figure.  In  this  proc- 
ess, however,  a  certain  definite  amount  of  water  vapor 


58     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

(depending  upon  the  temperature)  passes  into  the  tube  along 
with  the  oxygen,  so  that  the  observed  volume  of  the  gas 
in  the  tube  is  the  sum  of  the  volumes  of  the  oxygen  and 
of  the  water  vapor. 

Correction  for  aqueous  vapor  pressure.  The  volume 
which  the  pure  oxygen  will  occupy  under  any  definite 
conditions  of  temperature  and  pressure  may  be  determined 
as  follows: 

The  tube  A  (Fig.  29)  is  first  raised  or  lowered  until 
the  level  of  the  water  within  and  without  the  tube  is 
the  same.  Under  these  conditions  the  pressure  exerted 
by  the  atmosphere  upon  the  surface  of  the  water  in  the 
cylinder,  and  tending  to  force  the  liquid  up  into  the  tube  A, 
is  in  equilibrium  with  the  sum  of  the  pressures  exerted 
by  the  water  vapor  and  the  oxygen  within  the  tube  A  and 
tending  to  prevent  the  water  from  rising  in  the  tube.  If 
we  imagine  the  water  vapor  to  be  suddenly  removed,  but 
the  volume  to  be  kept  unchanged,  it  is  evident  that  the 
oxygen  would  occupy  the  entire  volume,  and  that  it  would 
exert  a  pressure  equal  to  the  atmospheric  pressure  less  the 
pressure  that  had  been  exerted  by  the  water  vapor  within 
the  tube.  If  we  know  the  latter  pressure,  we  can  subtract 
it  from  the  atmospheric  pressure  and  so  obtain  the  pres- 
sure exerted  by  the  oxygen  alone  at  the  given  temperature. 
The  pressure  of  water  vapor  for  all  ordinary  tempera- 
tures has  been  determined,  and  a  table  of  such  pressures 
is  given  in  the  Appendix.  The  pressures  are  expressed 
for  convenience  in  terms  of  the  height  of  a  column  of 
mercury  which  the  pressure  will  sustain. 

Example  of  correction  for  the  pressure  of  water  vapor. 
We  are  now  in  a  position  to  make  correction  for  the 
pressure  of  water  vapor,  and  an  example  will  make  the 
procedure  clear: 


THE  GAS  LAWS;  THE  KINETIC  THEOEY      59 


Let  us  suppose  that  a  gas  collected  over  water  measures 
300  cc.  under  a  pressure  of  740  mm.  (as  indicated  by  the  baro- 
metric reading)  and  a  temperature  of  25°  (298°  absolute  scale). 
Calculate  the  volume  which  the  pure  gas  will  occupy  under 
standard  conditions  (0°  and  760  mm.  pressure). 

First  find  the  change  in  volume  due  to  change  in  pressure. 
The  gas  collected  over  water  is  subjected  to  a  pressure  of 
740  mm.  less  the  pressure  exerted  by  the 
aqueous  vapor  at  25°;  namely,  23.69mm.; 
740  -  23.69  =  716.31.  Now  apply  the  ~y~  -  -A 

law  of  Boyle : 

300  x  716.31  =  760  x  V ;  or  V=  282.7  cc. 

Next  make  the  correction  for  temper- 
ature: 

— ^—  =  — —  ;  or  V  =  258.9  cc. 


FIG.  30.      Method      of 

measuring    the    vapor 

pressure  of  a  liquid 


Determination  of  vapor  pressure.  The 
pressure  of  the  vapor  of  a  liquid  at 
any  temperature  may  be  determined 
experimentally  in  the  following  way: 
Two  long  barometer  tubes  are  filled 
with  mercury  and  inverted  in  an  open 
vessel  of  the  same  liquid  (Fig.  30).  A 
few  drops  of  the  liquid  to  be  examined 
are  introduced  under  the  open  end  of 
one  of  the  tubes,  the  liquid  so  intro- 
duced immediately  rising  to  the  top  of  the  mercury  column. 
Evaporation  at  once  takes  place  and,  because  of  the  pressure 
of  the  gas  so  formed,  the  mercury  column  falls  to  some  extent. 
When  equilibrium  is  reached,  the  difference  in  level  of  the 
mercury  in  the  two  tubes,  included  between  the  dotted  lines 
A  and  B  in  the  figure,  will  correspond  to  the  pressure  of  the 
vapor  of  the  liquid  expressed  in  millimeters  of  mercury.  The 
tubes  may  be  surrounded  by  jackets  through  which  heated 
liquids  are  circulated,  so  that  any  desired  temperature  may 
be  secured. 


60     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  meaning  of  laws  in  science.  The  law  of  Boyle  and 
the  law  of  Gay-Lussac  are  merely  general  statements  in 
regard  to  the  conduct  of  gases  as  determined  by  experi- 
ment. Like  all  other  natural  laws,  they  offer  no  explana- 
tion of  the  facts  which  they  state,  nor  do  they  place  any 
restriction  upon  nature  which  compels  obedience,  as  the 
laws  enacted  by  a  legislature  bind  society.  They  are 
simply  concise  statements  of  what  might  be  called  the 
habits  of  nature  as  observed  in  experiment. 

Forming  a  theory.  It  is  certainly  a  very  striking  fact 
that  all  gaseous  substances  behave  in  so  simple  a  manner, 
quite  irrespective  of  their  chemical  nature.  It  would 
appear  most  probable  that  this  must  be  due  to  some  very 
simple  mechanical  structure  which  all  gases  have  in  com- 
mon, and  the  mind  at  once  begins  to  imagine  a  mechanical 
model  which,  if  real,  would  act  in  the  same  manner.  The 
process  of  constructing  a  mental  picture  of  this  kind  is 
called  forming  a  theory.  The  theory  which  has  proved  to 
be  the  most  satisfactory  in  connection  with  the  properties 
of  gases  is  known  as  the  kinetic  theory.  It  should  be 
noticed  that  laws  are  deduced  from  experiment  and  ex- 
press the  undoubted  truth,  while  theories  are  products 
of  the  imagination  and  are  always  open  to  error. 

The  kinetic  theory.  This  theory  imagines  the  following 
general  assumptions  to  be  true : 

1.  All  gases  are  made  up  not  of  continuous  matter  but 
of  extremely  minute  particles  relatively  far  apart.  These 
particles  are  called  molecules.  Millikan  has  calculated  that 
1  cc.  of  a  gas  under  standard  conditions  contains  27.09 
billion  billion  molecules.  This  assumption  satisfactorily 
accounts  for  the  fact  that  gases  may  be  compressed, 
since  pressure  would  merely  crowd  the  molecules  into  a 
smaller  space. 


THE  GAS  LAWS;  THE  KINETIC  THEORY      61 

2.  The    molecules    are    constantly    moving    in    straight 
lines   with   enormous   velocities,   and   are  hitting   against 
each  other  and  the  sides  of  the  containing  vessel  and  re- 
bounding.    For  example,  the  molecules  of  hydrogen  gas 
move   at  a  velocity  of  about  a  mile   per  second.     This 
assumption    makes   it   clear   why   gases   tend    to   expand 
indefinitely;   also  .why  gases  exert  a  pressure  against  the 
sides  of  any  vessel  filled  with  the  gas. 

3.  The  effect  of  heat  upon  a  gas  is  to  increase  the  ve- 
locity of  the  molecules.    If  the  velocity  of  the  molecules 
is  increased,  the  gas  will  tend  to  expand,  and  this  satis- 
factorily accounts  for  the  relation  between  the  volume  of 
a  gas  and  the  temperature  as  stated  in  the  law  of  Gay- 
Lussac. 

4.  For  certain  reasons  that  will  be  brought  out   in   a 
later  chapter,  it  is  also  assumed  that  all  the  molecules  of 
any  given  gaseous  element  (or  compound)  have  the  same 
weight,  but  that  the  molecules  of  different  elements  have 
different   weights;    also  that  equal  volumes  of  all  gases 
under  the   same  conditions  of  temperature  and  pressure 
contain  the  same  number  of  molecules.    This  last  assump- 
tion was  advanced  by  the  Italian  physicist  Avogadro  and 
is  known  as  Avogadro 's  hypothesis  (p.  281). 

Value  of  a  theory.  The  value  of  such  a  theory  as  the 
kinetic  theory  is  at  once  apparent.  It  presents  a  mental 
picture  which  assists  the  memory  in  retaining  a  great 
variety  of  facts.  It  suggests  many  experiments  which 
otherwise  might  never  be  undertaken,  for  our  first  impulse, 
after  forming  such  a  theory,  is  to  test  it  experimentally 
in  every  possible  way.  It  enables  us  to  form  a  probable 
opinion  in  cases  where  experiment  has  not  yet  made  a 
definite  decision.  It  often  leads  to  the  detection  of  errors 
which  have  crept  into  the  body  of  our  knowledge  as  the 


62     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

result  of  faulty  experiments.  There  is,  however,  a  real 
peril  in  accepting  a  theory.  The  whole  picture  may  be 
wrong,  yet  it  may  seem  so  plausible  that  we  rest  con- 
tented with  it  and  fail  to  see  its  faults.  Its  very  plau- 
sibility may  prevent  us  from  making  experiments  which 
would  disclose  the  error  in  the  theory  and  put  us  011  the 
right  track. 

All  such  theories  are  best  regarded  as  mere  conveniences. 
Doubtless  they  express  the  true  nature  of  things  in  many 
cases,  but  in  others  they  do  not.  They,  are  useful  so  long 
as  we  regard  them  as  conveniences,  and  open  to  constant 
revision  and  modification  or  to  rejection  as  our  knowledge 
grows ;  but  they  are  a  real  disadvantage  when  we  come 
to  regard  them  as  the  final  and  unchangeable  truth. 

EXERCISES 

1.  Why  does  a  balloon  tend  to  fall  at  night  and  rise  at  midday? 

2.  Why  is  the  bottom  of  a  balloon  left  open  and  not  tightly  closed? 

3.  How  can  you  change  the  readings  on  the  Fahrenheit  scale 
into  readings  on  the  centigrade  scale? 

4.  What  evidence  can  you  give  tending  to  show  that  the  amount 
of  water  vapor  taken  up  by  the    atmosphere  increases  with  the 
temperature  ? 

5.  Why  does  the  carburetor  on  a  motor  car  have  to  be  adjusted 
with  changes  in  the  atmospheric  conditions,  in  order  to  secure  the 
greatest  efficiency? 

6.  A  gas   measures    200  cc.  when    subjected   to    a   pressure   of 
740  mm.    What  volume  will  it  occupy  if  the  pressure  is  increased 
to  760  mm.,  the  temperature  remaining  constant  ? 

7.  A  gas  measured  150  cc.  in  a  laboratory  in  which  the  tempera- 
ture was  18°.    What  volume  would  the  gas  occupy  if  the  temperature 
is  reduced  to  0°,  the  pressure  remaining  constant? 

8.  A  certain  volume  of  gas  is  measured  when  the  temperature 
is  20°  and  the  pressure  740  mm.   (a)  At  what  temperature  would  its 


THE  GAS  LAWS;  THE  KINETIC  THEORY      63 

volume  be  doubled,  the  pressure  remaining  constant?  (i)  At  what 
pressure  would  its  volume  be  doubled,  the  temperature  remaining 
constant  ? 

9.  A  certain  weight  of  hydrogen,  subjected  to  atmospheric 
pressure,  occupied  a  volume  of  1250  cc.  in  a  laboratory  in  which 
the  readings  of  the  thermometer  and  barometer  were  18°  and  746  mm. 
respectively,  (a)  What  volume  would  this  gas  occupy  under  stand- 
ard conditions  ?  (6)  Calculate  the  weight  of  this  volume  of  hydrogen. 

10.  A  gas  subjected  to  atmospheric  pressure  occupied  a  volume 
of  1000  cc.  in  a  laboratory  at  20°  and  740  mm.   The  next  day  the 
temperature  of  the  laboratory  fell  to  12°,  while  the  barometric  pres- 
sure increased  to  752  mm.    What  volume  did  the  gas  occupy  under 
these  conditions? 

11.  (a)  2500  cc.  of   oxygen   measured   over  water  at  20°  and 
740mm.  would  occupy  what  volume  under  standard  conditions? 
(ft)  Calculate  the  weight  of  this  volume  of  the  gas.    (c)  What  weight 
of  potassium  chlorate  would  be  required  to  prepare  this  amount 
of  oxygen? 

12.  I  wish  to  prepare  100 1.  of  hydrogen,  measured  over  water  in  a 
laboratory  where  the  temperature  and  pressure  are  respectively  16° 
and  750  nun.    What  weight  of  hydrogen  sulfate  would  be  required? 

13.  Oxygen  was  compressed  in  a  50-liter  cylinder  at  20°  until  the 
pressure  was  equal  to  100  atmospheres.    What  volume  would  this 
occupy  under  a  pressure  of  1  atmosphere,  the  temperature  remaining 
constant? 

14.  Suppose  that  the  cylinder  of  oxygen  referred  to  in  problem  13 
were  heated  to  500°  during  a  fire.    What  pressure  would  the  gas 
exert  at  this  temperature? 


CHAPTER  VI 
WATER;  HYDROGEN  PEROXIDE 

Historical.  Following  the  discovery  of  hydrogen,  Caven- 
dish made  a  careful  study  of  the  properties  of  the  gas.  In 
the  course  of  his  experiments  he  exploded  a  mixture  of 
hydrogen  and  oxygen  and  observed  that  a  small  amount 
of  a  dewlike  substance  was  formed.  He  was  able  to  obtain 
a  sufficient  amount  of  the  liquid  to  make  a  study  of  its 
properties,  and  showed  that  it  was  pure  water.  Cavendish 
did  not  perceive  the  full  meaning  of  his  discovery,  how- 
ever, and  it  remained  for  Lavoisier,  a  few  years  later,  to 
repeat  and  properly  interpret  the  experiments  of  Caven- 
dish. He  proved  beyond  doubt  that  the  water  which 
Cavendish  had  obtained  resulted  from  the  union  of  the 
hydrogen  and  oxygen,  and  that  water  must  be  regarded 
as  a  compound  of  these  two  elements. 

Occurrence.  The  great  abundance  and  wide  distribution 
of  water  are  facts  familiar  to  all.  Vast  areas  of  the  colder 
regions  of  the  globe  are  covered  with  it  in  the  form  of 
ice,  while  in  the  liquid  state  it  covers  about  five  sevenths 
of  the  earth's  surface,  reaching  in  some  places  a  depth  of 
nearly  six  miles.  Large  quantities  occur  in  the  soil,  and 
as  a  vapor  it  is  an  essential  constituent  of  the  atmosphere. 
It  likewise  constitutes  more  than  half  the  weight  of  living 
organisms.  For  example,  nearly  70  per  cent  of  the  human 
body  is  water.  The  water  content  of  some  of  the  more 
common  foods  is  given  in  the  table  on  page  344. 
64 


WATER;  HYDROGEN  PEROXIDE  65 

Composition  of  natural  waters.  Water  as  it  occurs  in 
nature  always  contains  more  or  less  matter  derived  from 
the  rocks  and  soils  with  which  it  comes  in  contact.  When 
such  water  is  evaporated,  this  matter  is  left  behind  in 
solid  form.  Even  rain  water,  which  is  the  purest  natural 
water,  contains  dust  particles  and  gases  dissolved  from  the 
atmosphere.  The  foreign  matter  in  natural  waters  is  of 
two  kinds ;  namely,  mineral  and  organic. 

1.  Mineral  matter.  The  amount  and  nature  of  the  min- 
eral substances  present  in  any  given  water  vary  with  the 
nature   of  the  rocks  and  soil  with  which  the  water  has 
been  in  contact.   The  mineral  substances  ordinarily  present 
in  fresh  waters  are  common  salt  and  compounds  of  cal- 
cium, magnesium,  and  iron.    Water  containing  any  con- 
siderable  amounts   of   mineral   matter   does   not   form   a 
lather  with  soap,  and  is  termed  hard  water',  or  if  a  large 
amount  of  mineral  matter  is  present,  it  is  called  a  mineral 
water.    Water  containing  little  or  no  mineral  matter,  such 
as  rain  water,  is  termed  soft  water.    One  liter  of  an  aver- 
age river  water  contains  about  0.1 75  g.  of  mineral  matter. 
The  water  of  the  ocean  contains  about  40  g.  of  mineral 
matter  to  the  liter,  more  than  three  fourths  of  which  is 
common   salt.    The  water   of   Great  Salt  Lake  contains 
from  150  to  200  g.  of  mineral  matter  in  1  1. 

2.  Organic  matter.   In  addition  to  mineral  matter  natural 
waters  contain  more  or  less  organic  matter  in  solution  or 
held  in  suspension.    This  consists  not  only  of  inanimate 
matter,  derived  from  the  decay  of  organic  bodies  on  the 
earth's  surface  or  present  in  sewage,  but  also  of  certain 
forms  of  living  microorganisms  which  usually  accompany 
such    products.     Waters    taken    from    shallow    wells    or 
streams  in  thickly  populated  districts  are  likely  to  contain 
a  considerable  quantity  of  such  matter. 


66     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Effect  upon  health  of  the  foreign  matter  in  water.  Since 
natural  waters  constitute  the  ordinary  supply  for  drinking 
and  household  purposes,  the  effect  of  the  foreign  matter  in 
such  waters  when  taken  into  the  system  is  a  question  of 
great  importance.  Experience  has  shown  that  the  mineral 
matter  commonly  found  in  water  is  not,  as  a  rule,  injuri- 
ous to  health.  In  fact,  the  presence  of  a  certain  amount 
of  such  matter  is  probably  advantageous,  supplying  a 
portion  of  the  mineral  constituents  necessary  for  the 
formation  of  the  solid  tissues  of  the  body. 

The  organic  matter  present  in  water  consists  of  inani- 
mate products  as  well  as  of  living  microorganisms.  The 
amount  of  the  former  commonly  present  in  a  water  used 
for  drinking  purposes  is  so  small  that  it  is  virtually  with- 
out effect  upon  health.  On  the  other  hand,  the  presence, 
in  water,  of  any  considerable  number  of  microorganisms 
renders  it  dangerous  as  a  drinking  water.  It  is  true  that 
many  of  these  organisms  are  without  injurious  effect,  but 
it  is  likewise  true  that  others  are  the  direct  cause  of 
disease.  It  is  known  that  a  transmissible  disease  such  as 
typhoid  fever  is  due  to  certain  microorganisms  which  find 
entrance  into  the  body.  It  is  easily  possible  for  these  organ- 
isms to  find  their  way,  through  sewage,  from  a  person  afflicted 
with  the  disease  into  a  poorly  protected  water  supply,  and  so 
contaminate  the  water.  It  is  largely  in  this  way  that  typhoid 
fever  »«  spread.  The  general  conclusion  may  therefore  be 
drawn  that,  save  in  exceptional  cases,  any  sickness  trace- 
able to  the  water  supply  is  due  to  the  presence  in  the 
water,  not  of  mineral  matter,  or  even  of  inanimate  organic 
matter,  but  to  certain  living  microorganisms. 

The  detection  of  impurities  in  water.  The  total  amount 
of  solid  matter  present  in  any  given  water  is  easily  deter- 
mined by  evaporating  a  definite  volume  of  the  water  to 


WATER;  HYDROGEN  PEROXIDE  67 

dryness  and  weighing  the  residue.  This  residue  may  then 
be  subjected  to  further  investigation  and  the  nature  of 
the  mineral  matter  determined.  An  examination  of  this 
kind  is  called  a  mineral  analysis.  Such  an  analysis  is  of 
importance  in  determining  whether  or  not  a  water  is 
adapted  for  manufacturing  purposes,  such  as  for  use  in  a 
steam  boiler. 

On  the  other  hand,  if  one  wishes  to  determine  whether 
a  water  is  wholesome  for  drinking,  a  so-called  sanitary 
analysis  is  required.  Such  an  analysis  includes  not  only 
the  determination  of  the  organic  matter  present  in  the 
water  but  also  of  the  decomposition  products  formed 
by  the  decay  of  such  matter  (chiefly  ammonia,  nitrites, 
and  nitrates).  From  what  has  been  said  it  might  be  in- 
ferred that  a  bacteriological  examination  alone  would  be 
sufficient.  While  it  is  true  that  such  an  examination  is 
of  the  greatest  importance,  it  is  equally  true  that  the 
determination  of  the  inanimate  organic  matter  present, 
together  with  the  products  of  its  decomposition,  is  of 
great  value  and  supplements  the  knowledge  gained  from 
a  bacteriological  examination ;  for  the  disease-producing 
organisms  find  their  way  into  a  water  supply  through  the 
sewage  or  drains,  and  are  therefore  accompanied  by  other 
organic  matter,  the  presence  of  which  hi  a  water  supply 
at  once  indicates  pollution.  Such  a  water  should  therefore 
not  be  used,  for,  although  it  may  be  temporarily  free  from 
disease-producing  organisms,  the  conditions  are  such  that 
their  introduction  may  take  place  at  any  time. 

It  may  be  added  that  the  physical  properties  of  a 
drinking  water  rarely  give  any  conclusive  evidence  as  to 
its  purity.  A  water  may  be  unfit  for  drinking  and  yet 
be  perfectly  clear  and  odorless.  Neither  can  any  reliance 
be  placed  on  the  simple  methods  sometimes  given  for 


68     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

testing  the  purity  of  water.  Only  the  trained  chemist  and 
bacteriologist  can  carry  out  such  methods  of  analysis  as 
are  trustworthy. 

Purification  of  water.  Three  general  methods  are  used 
for  the  purification  of  water ;  namely,  distillation,  boiling, 
and  filtration. 

1.  Distillation.  The  most  effective  way  of  purifying  ordi- 
nary water  is  by  the  process  of  distillation.  This  consists 
in  boiling  the  water  and  condensing  the  resulting  steam. 

In  the  laboratory 
the  process  is  usu- 
ally conducted  as 
follows : 

Ordinary    water 
is  poured  into  the 
flask   A     (Fig.  31) 
and     boiled.      The 
steam  is  conducted 
through    the    con- 
denser B,  commonly 
FIG.  31.   The  distillation  of  water  as  carried        kn<>wn  as  a  Liebig 
on  in  the  laboratory  condenser,      which 

consists  essentially 

of  a  narrow  glass  tube  sealed  within  a  larger  one,  the  space 
between  the  two  being  filled  with  cold  water,  which  enters  at  C 
and  escapes  at  D.  In  this  way  the  inner  tube  is  kept  cool  and 
the  steam  in  passing  through  it  is  condensed.  The  water  formed 
by  the  condensation  of  the  steam  collects  in  the  receiver  E. 

The  water  formed  by  the  condensation  of  .steam  is 
known  as  distilled  water.  The  mineral  matter  present  in 
the  original  water  is  not  volatile  and'  remains  in  the  con- 
tainer in  which  the  water  is  boiled.  The  organic  matter 
is  also  largely  left  in  the  container.  A  small  amount  of 
it,  however,  may  be  decomposed  into  volatile  products,  in 


WATER;  HYDROGEN  PEROXIDE 


69 


which  case  these  will  pass  over  with  the  steam  and  be 
present  in  the  distilled  water.  The  percentage  of  such 
matter  in  distilled  water  is  so  small,  however,  that  it  is 
without  effect  in  most  of  the  chemical  processes  in  which 
pure  water  is  employed;  in  a  very  few  cases  where  ex- 
treme purity  is  required,  further  treatment  is  necessary. 


FIG.  32.   The  distillation  of  water  for  commercial  purposes 

Distilled  water  is  used  by  the  chemist  in  almost  all  of 
his  work.  Large  quantities  are  also  used  in  the  manufac- 
ture of  ice,  as  well  as  for  drinking. 

Commercial  distillation.  In  preparing  distilled  water  on  a 
large  scale  the  steam  is  generated  in  a  metal  boiler  A  (Fig.  32) 
and  is  conducted  through  the  pipe  B  to  the  condensing  coil  C, 
made  of  tin.  This  pipe  is  wound  into  a  spiral  and  is  sur- 
rounded by  cold  water,  which  enters  at  D  and  flows  out  at  E. 
The  distilled  water  is  collected  in  a  suitable  container  F. 


70     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

2.  Boiling.    In  purifying  water  for  drinking  purposes  it 
is  only  necessary  to  remove  or  destroy  the  microorganisms 
present.    When   the   amount  of  water   to   be   purified  is 
small,  as  is  the  case  with  the  household  supply  for  drink- 
ing, this  is  most  conveniently  accomplished  by  boiling  the 
water  for  ten   or  fifteen  minutes.    While  the    organisms 
are  destroyed  in  a  short  time  by  moist  heat,  even  severe 
cold  has  been  found  to  have  comparatively  little   effect 
upon  them. 

3.  Filtration.    On  a  small  scale,  water  is  filtered  in  two 
ways:    (1)  by  passing  it  through  some  porous  material, 
such  as  charcoal,  or  (2)  by  forcing  it  through  porous  clay 
ware,  as  is  done  in  the  Chamberlain-Pasteur  filter.    While 
such  filters,  if  kept  clean  and  in  good  condition,  remove 
most  of  the  organic  matter,  they  do  not  remove  mineral 
matter  except  such  as  is  held  in  suspension.    These  house- 
hold filters  are  not  easily  kept  in  order  and  soon  become 
ineffective.    They  are  no  longer  used  to  any  great  extent. 

City  filtration.  Many  cities  find  it  necessary  to  take  their 
water  supply  from  rivers.  The  rivers  in  thickly  populated 
districts  are  almost  certain  to  be  contaminated  with  organic 
matter,  suggesting  the  possible  presence  of  disease  germs. 
Such  water  is  a  constant  menace  to  the  health  of  the  city,  so 
that  it  is  of  the  greatest  importance  to  find  some  way  of  puri- 
fying it  effectively  on  a  large  scale.  This  is  done  by  filtration. 
Two  general  kinds  of  filters  are  in  use : 

1.  Slow  sand  filters.  These  consist  of  large  beds  of  sand  and 
gravel,  through  which  the  water  passes  slowly.  Some  of  the 
impurities  are  strained  out,  while  others  are  decomposed  by 
the  action  of  certain  kinds  of  microorganisms  which  collect 
in  a  jellylike  layer  on  the  surface  of  the  filter.  The  purified 
water  passes  into  a  porous  pipe  from  which  it  is  pumped  into 
the  city  mains.  The  filters  are  covered  to  protect  the  water 
and  prevent  it  from  freezing. 


WATER;  HYDROGEN  PEROXIDE 


71 


2.  Mechanical  fillers.  In  these  the  water,  before  filtration,  is 
run  into  large  tanks  and  mixed  intimately  with  certain  com- 
pounds, such  as  aluminium  sulfate  or  iron  sulfate,  which  form 
in  the  water  a  small  amount  of  gelatinous  solid.  The  water  is 
then  run  into  settling  basins,  where  the  gelatinous  solid  slowly 
settles  to  the  bottom,  carrying  with  it  much  of  the  organic 
matter  present.  The  partially  clarified  water  is  then  filtered 
through  sand  and  gravel  (Fig.  33).  When  this  method  is  em- 
ployed it  is  customary  to  add  some  germicide,  such  as  chlorine 
(Fig.  77),  to  the  water 
in  order  to  destroy 
any  microorganisms 
not  removed  by  the 
gelatinous  solid.  The. 
mechanical  filter  is 
largely  replacing  the 
slow  sand  filter,  since 
the  process  is  much 
more  rapid  and  just 
as  effective. 

The  effect  of  the 
filtration  of  the  water 
supply  upon  the  health 
of  a  city  is  shown  by 
the  fact  that  in  gen- 
eral the  number  of  cases  of  typhoid  fever  in  cities  which  have 
introduced  an  effective  water-purification  system  has  been  de- 
creased by  about  75  per  cent. 

Self-purification  of  water..  It  has  long  been  known  that  water 
contaminated  with  organic  matter  tends  to  purify  itself  when 
exposed  to  the  air  (p.  34).  This  is  due  to  the  fact  that  air  is 
somewhat  soluble  in  water,  and  that  the  dissolved  oxygen,  in 
the  presence  of  certain  microorganisms,  gradually  oxidizes  the 
organic  matter  present  in  the  water ;  when  this  is  destroyed, 
the  organisms  present  die  for  lack  of  food.  This  process,  how- 
ever, cannot  be  relied  upon  to  purify  a  contaminated  water  so 
as  to  render  it  safe  for  drinking  purposes. 


FIG.  33.    Sand  and  gravel  filter  in  a  modern 
city  filtration  plant 


72     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Properties  of  water.  Pure  water  is  an  odorless  and 
tasteless  liquid,  colorless  in  thin  layers  but  having  a 
bluish  tinge  when  observed  through  a  considerable  thick- 
ness. It  solidifies  at  0°  and  boils  at  100°  under  the  nor- 
mal pressure  of  1  atmosphere.  When  water  is  cooled,  it 
steadily  contracts  until  the  temperature  of  4°  is  reached ; 
at  lower  temperatures  it  expands.  Water  is  remarkable 
for  its  ability  to  dissolve  other  substances,  and  is  the 
most  general  solvent  known.  Even  such  substances  as 
glass  and  various  kinds  of  rocks  which  are  ordinarily 
regarded  as  insoluble  dissolve  in  water,  but  to  a  very 
limited  extent.  Chemists  usually  employ  aqueous  solu- 
tions of  substances  rather  than  the  substances  themselves, 
since,  as  a  rule,  chemical  action  takes  place  more  readily 
in  solution. 

Chemical  conduct.  Water  is  a  very  stable  substance ;  in 
other  words,  it  does  not  undergo  decomposition  readily. 
To  decompose  it  into  its  elements  by  heat  alone  requires 
a  very  high  temperature.  Even  at  2500°  only  about  10  per 
cent  of  the  water  heated  is  decomposed.  Though  very  stable 
toward  heat,  water  can  be  decomposed  in  other  ways,  as 
by  the  action  of  the  electric  current  or  by  certain  metals. 

Though  containing  88.81  per  cent  of  oxygen,  water  is 
not  a  good  oxidizing  agent,  because  of  its  great  stability. 
However,  certain  metals,  as  well  as  carbon,  can  be  oxidized 
by  very  hot  steam,  the  hydrogen  being  set  free.  Water 
combines  directly  with  many  compounds,  forming  substances 
called  hydrates  (p.  251).  Blue  vitriol  and  alum  are  good 
examples  of  such  hydrates. 

Heat  of  formation  and  heat  of  decomposition  are  equal.  The 
fact  that  a  very  high  temperature  is  necessary  to  decompose 
water  into  hydrogen  and  oxygen  is  in  accord  with  the  fact  that 
a  great  deal  of  heat  is  evolved  by  the  union  of  hydrogen  and 


WATER;  HYDROGEN  PEROXIDE  73 

oxygen  (p.  49),  for  it  lias  been  proved  that  the  heat  necessary 
to  decompose  a  compound  into  its  elements  (heat  of  decom- 
position) is  equal  to  the  heat  evolved  in  the  formation  of  the 
same  compound  from  its  elements  (heat  of  formation). 

The  determination  of  the  exact  composition  of  water. 
Many  very  careful  experiments  have  been  made  for  the 
purpose  of  determining,  with  as  great  accuracy  as  possible, 
the  ratio  in  which  hydrogen  and  oxygen  are  present  in 
water,  and  it  is  worth  our  while  to  study  somewhat  in 
detail  the  methods  which  have  been  employed,  since  they 
serve  to  illustrate  in  a  general  way  the  methods  used  in 
determining  the  composition  of  other  compounds. 

Two  general  methods  of  procedure  are  available  for 
determining  the  composition  of  a  compound :  first,  the 
method  of  analysis,  in  which  a  given  weight  of  the  com- 
pound is  separated  either  directly  or  indirectly  into  its 
constituent  elements  and  the  identity  and  weight  of  each 
determined ;  second,  the  method  of  synthesis,  which  consists 
in  determining  the  proportion  in  which  the  constituent 
elements  unite  to  form  the  compound,  and  which  is 
therefore  just  the  opposite  of  analysis. 

1.  Methods  based  on  analysis.  It  will  be  recalled  that 
water  may  be  easily  decomposed  into  its  constituents  by 
the  electric  current.  It  would  seem  probable  that  the 
exact  composition  of  water  could  easily  be  determined  in 
this  way,  since  the  volumes  of  the  gases  liberated  can 
readily  be  measured  with  accuracy,  and  if  we  know  their 
densities,  we  can  calculate  their  weights.  When  the  ex- 
periment is  carried  out,  however,  the  results  obtained  are 
not  concordant,  although  in  general  the  volume  of  the 
hydrogen  liberated  is  slightly  more  than  double  the  vol- 
ume of  the  oxygen.  Experiments  prove  that  the  method 
is  subject  to  several  sources  of  error.  For  example,  the 


74     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


water  through  which  the  liberated  gases  bubble  dissolves 
more  of  the  oxygen  than  of  the  hydrogen.  The  ratio 
between  the  amounts  of  hydrogen  and  oxygen  obtained  in 
this  process,  therefore,  does  not  represent  with  great  accu- 
racy the  ratio  in  which  they  are  combined  in  water.  More 
accurate  results  are  obtained  by  the  synthetic  methods  de- 
scribed in  the  succeeding  paragraphs. 

2.  Methods  based  on  synthesis.  In  the 
synthetic  methods  we  determine  the 
quantities  of  oxygen  and  hydrogen 
which  combine  directly  to  form  water. 
A  description  of  the  method  as  carried 
out  for  purposes  of  illustration  in  the 
lecture  room  will  serve  to  show  the 
general  principle  involved. 

The  combination  of  the  two  gases  is 
brought  about  in  a  tube  called  a  eudiome- 
ter. This  is  a  graduated  glass  tube  about 

GO  cm.  long  and  2  cm.  wide,  closed  at  one 

FIG.  34.      A     simple  ,    /T,.      n4\      -\T         i-u        i        ^         i    x. 

form  of   eudiometer     end  (Fig.  34).    Near  the  closed  end  two 

platinum  wires  are  fused  through  the  glass, 

the  ends  of  the  wires  within  the  tube  being  separated  by  a 
space  of  2  or  3  mm.  The  tube  is  entirely  filled  with  mercury 
and  inverted  in  a  vessel  of  the  same  liquid.  Pure  hydrogen 
is  passed  into  the  tube  until  it  is  about  one  fourth  filled. 
The  tube  is  then  lowered  until  the  mercury  stands  at  the  same 
level  inside  and  outside  the  tube,  and  the  reading  of  the  volume 
of  the  hydrogen  is  taken.  Approximately  an  equal  volume  of 
pure  oxygen  is  then  introduced,  and  the  volume  is  again  taken. 
This  gives  the  total  volume  of  the  two  gases.  From  this  the 
volume  of  the  oxygen  introduced  may  be  determined  by 
subtracting  from  it  the  volume  of  the  hydrogen. 

The  combination  of  the  two  gases  is  now  brought  about  by 
connecting  the  two  platinum  wires  with  an  induction  coil  and 
passing  a  spark  from  one  wire  to  the  other.  Immediately  a 


WATER;  HYDROGEN  PEROXIDE  75 

slight  explosion  occurs.  The  mercury  in  the  tube  is  at  first 
depressed,  because  of  the  expansion  of  the  gases  due  to  the 
heat  generated,  but  it  at  once  rebounds,  taking  the  place  of 
the  gases  which  have  combined  to  form  water.  The  volume 
of  the  water  in  the  liquid  state  is  so  small  that  it  may  be 
disregarded  in  the  calculations. 

In  order  that  the  temperature  of  the  residual  gas  and  the 
mercury  may  become  uniform,  the  apparatus  is  allowed  to  stand 
for  a  few  minutes,  and  the  volume  of  the  gas  is  taken.  The 
residual  gas  is  then  tested  in  order  to  ascertain  whether  it  is 
hydrogen  or  oxygen,  since  experiments  have  proved  that  it  is 
never  a  mixture  of  the  two.  From  the  information  thus  obtained 
the  composition  of  the  water  may  be  calculated. 

Calculation  of  composition.  Thus,  suppose  the  readings  were 
as  follows : 

Volume  of  hydrogen 20.3  cc. 

Volume  of  hydrogen  and  oxygen 38.7  cc. 

Volume  of  oxygen 18.4  cc. 

Volume  of  gas  left  after  combination  has  taken 

place  (found  to  be  oxygen) 8.3  cc. 

We  have  thus  found  that  20.3  cc.  of  hydrogen  have  com- 
bined with  18.4  cc.  minus  8.3 cc.  (or  10.1  cc.)  of  oxygen;  or 
approximately  2  volumes  of  hydrogen  have  combined  with 
1  volume  of  oxygen.  Since  oxygen  is  15.9  times  as  heavy  as 
hydrogen,  the  proportion  by  weight  in  which  the  two  gases 
combine  is  1  part  of  hydrogen  to  7.94  parts  of  oxygen. 

A  convenient  form  of  eudiometer.  A  form  of  eudiometer  rep- 
resented in  Fig.  35,  and  different  from  that  shown  in  Fig.  34, 
is  sometimes  used,  since  it  is  easier  by  means  of  this  to  obtain 
the  gases  under  the  same  conditions  of  temperature  and  pres- 
sure in  order  to  make  comparisons.  With  this  apparatus  it  is 
easily  possible  to  take  the  readings  of  the  volumes  under  the 
same  conditions  of  temperature  and  pressure,  and  thus  compare 
them  directly.  The  apparatus  is  filled  with  mercury  and  the 
gases  introduced  into  the  tube  A.  The  experiment  is  carried 
out  like  the  preceding  one,  except  that,  before  taking  the 


76     AX  ELEMENTARY  STUDY  OF  CHEMISTRY 


reading  of  the  gas  volumes,  mercury  is  either  added  to  the  tube 
B  or  withdrawn  from  it  by  means  of  the  stopcock  C,  until  it 
stands  at  exactly  the  same  height  in  both  tubes.  The  gas 
inclosed  in  tube  A  is  then  under  atmospheric  pressure.  The 
temperature  of  the  gas,  as  well  as  the  pressure  to  which  it 
is  subjected,  being  the  same  at  the  conclusion  of  the  experi- 
ment as  at  the  beginning,  the  volumes  of  the  hydrogen  and 
oxygen  and  of  the  residual  gas  may 
be  directly  compared  as  read  off  fron? 
the  tube. 


Method  used  by  Berzelius  and 
Dumas.  The  work  of  Berzelius  and 
Dumas  is  of  interest  from  a  histori- 
cal standpoint,  since  they  were  the 
first  to  determine  the  composition  of 
water  with  any  great  accuracy.  The 
method  used  is  a  very  ingenious 
one,  the  weights  of  the  hydrogen 
and  oxygen  being  determined  by 
indirect  methods  and  not  by  direct 
weighing  of  the  gases,  which  is  not 
easily  done.  The  method  was  first 
used  by  Berzelius  in  1820,  and  later, 
in  1843,  with  greater  refinement, 
by  Dumas. 


FIG.  35.    A  convenient 
form  of  eudiometer 


Details  of  the  experiment.  Fig.  36  illustrates  the  essential 
parts  of  the  apparatus  used  in  making  the  determination.  The 
glass  tube  B  contains  copper  oxide,  while  the  tubes  C  and  D 
are  filled  with  calcium  chloride,  a  substance  which  has  great 
affinity  for  water.  The  tubes  B  and  C,  including  their  con- 
tents, are  carefully  weighed,  and  the  apparatus  is  connected  as 
shown  in  the  figure.  A  slow  current  of  pure  hydrogen  is  then 
passed  through  A,  and  that  part  of  the  tube  B  which  contains 
copper  oxide  is  carefully  heated.  The  hydrogen  combines  with 


WATER;  HYDROGEN  PEROXIDE  11 

the  oxygen  of  the  copper  oxide  to  form  water,  which  is  ab- 
sorbed by  the  calcium  chloride  in  tube  C.  The  calcium  chloride 
in  tube  D  prevents  any  moisture  entering  tube  C  from  the  air. 
The  operation  is  continued  until  an  appreciable  amount  of 
water  has  been  formed.  The  tubes  B  and  C  are  then  weighed 
once  more.  The  loss  of  weight  in  the  tube  B  will  exactly  equal 
the  weight  of  oxygen  taken  up  from  the  copper  oxide  in  the 
formation  of  the  water.  The 'gain  in  weight  in  the  tube  C 
will  exactly  equal  the  weight  of  the  water  formed.  The  differ- 
ence in  these  weights  will  of  course  equal  the  weight  of  the 
hydrogen  present  in  the  water  formed  during  the  experiment. 


FIG.  36.    Apparatus  for  determining  the  ratio  by  weight  in  which  oxygen 
and  hydrogen  combine 

Dumas's  results.  The  results  secured  by  Dumas  in  1843 
may  be  summed  up  as  follows: 

Weight  of  water  formed 945.439  g. 

Oxygen  given  up  by  the  copper  oxide      .     .     .     840.161  g. 
Weight  of  hydrogen  present  in  water      .     .     .     105.278  g. 

According   to  this   experiment  the  ratio  of  hydrogen  to 
oxygen  in  water  is  105.278:840.161,  or  1:7.98. 

Morley's  results.  In  recent  years  the  American  chemist 
Morley  has  determined  the  composition  of  water  with  great 
care.  Extreme  precautions  were  taken,  and  the  hydrogen 
and  oxygen  which  combined,  as  well  as  the  water  formed, 
were  all  accurately  weighed.  According  to  Morley's  re- 
sults, 1  part  by  weight  of  hydrogen  combines  with  7.94 
parts  by  weight  of  oxygen  to  form  water. 


78     AN  ELEMENTARY  STUDY  OF  CHEMISTEY 

Comparison  of  results  obtained.  From  what  has  been 
described  it  is  easy  to  see  that  it  is  by  experiment  alone 
that  the  composition  of  a  compound  can  be  determined. 
Different  methods  may  lead  to  slightly  different  results. 
The  more  accurate  the  method  chosen,  and  the  greater 
the  skill  with  which  the  experiment  is  carried  out,  the 
more  accurate  will  be  the  results.  It  is  generally  con- 
ceded by  chemists  that  the  results  obtained  by  Morley  in 
reference  to  the  composition  of  water  are  the  most  accu- 
rate ones.  In  accordance  with  these  results,  then,  water 
must  be  regarded  as  a  compound  containing  hydrogen 
and  oxygen  in  the  ratio  of  1  part,  by  weight  of  hydrogen 
to  7.94  parts  by  weight  of  oxygen. 

Relation  between  any  given  volume  of  water  vapor  and 
the  volumes  of  the  hydrogen  and  oxygen  which  combine  to 
form  it.  When  the  quantitative  synthesis  of  water  is  car- 
ried out  at  ordinary  temperatures,  the  water  vapor  formed 
by  the  union  of  the  hydrogen  and  ox}rgen  at  once  con- 
denses. The  volume  of  the  resulting  liquid  is  so  small 
that  it  may  be  disregarded  in  making  the  calculations. 
If,  however,  the  experiment  is  carried  out 'at  a  tempera- 
ture of  100°  or  above,  the  water  vapor  formed  is  not 
condensed,  and  it  then  becomes  possible  to  compare  the 
volume  of  the  vapor  with  the  volumes  of  hydrogen  and 
oxygen  which  combined  to  form  it.  In  this  way  it  has 
been  proved  that  2  volumes  of  hydrogen  and  1  volume 
of  oxygen  combine  to  form  exactly  2  volumes  of  water 
vapor,  the  volumes  all  being  measured  under  the  same 
conditions  of  temperature  and  pressure.  It  will  be  noted 
that  the  relation  between  these  volumes  may  be  expressed 
by  whole  numbers.  It  will  be  found  from  subsequent  dis- 
cussions that  a  similar  statement  holds  in  reference  to 
all  gaseous  elements  which  combine  with  each  other. 


WATER;  HYDROGEN  PEROXIDE 


79 


The  form  of  apparatus  used  in  determining  the  relation 
between  the  volumes  of  hydrogen  and  oxygen  uniting  and  that 
of  the  aqueous  vapor  formed  is  illustrated  in  Fig.  37.  The 
arm  A  of  the  eudiometer  in  which  the  combination  of  the 
gases  is  effected  is  surrounded  by  a  tube  through  which  is 
passed  steam  or,  preferably,  the  vapor  of  some  liquid  boil- 
ing above  100°  (ainyl  alcohol  is  often  used).  A  mixture  of 
2  volumes  of 
hydrogen  with 
1  of  oxygen  is 
introduced  into 
the  eudiometer. 
A  suitable  liq- 
uid is  then 
boiled  in  the 
flask  B.  The 
resulting  vapor 
is  conducted 
through  the 
space  between 
the  tube  A  and 


FIG.  37.    Apparatus  for  determining  the   ratio  by 

volume  in  which  oxygen  and  hydrogen  combine,  and 

the  ratio  of  each  of  these  volumes  to  the  volume  of 

the  steam  formed 


the  outer  tube 
and  is  then  con- 
densed as  shown 
in  the  figure. 
When  the  vol- 
ume of  the  mixed  gases  in  A  has  become  stationary,  showing 
that  the  temperature  of  the  gases  is  the  same  as  that  of  the 
vapor,  and  the  pressure  adjusted  as  in  the  former  experiment 
(see  Fig.  35  and  description),  the  reading  on  the  eudiometer 
tube  is  noted.  The  union  of  the  two  gases  is  then  effected  by 
an  electric  spark  from  an  induction  coil  C.  After  the  union 
has  taken  place,  the  pressure  is  adjusted,  and  the  reading  again 
noted  after  the  volume  of  the  vapor  has  become  constant.  The 
volume  of  the  vapor  thus  obtained  can  be  compared  directly 
with  the  volumes  of  the  hydrogen  and  oxygen  which  united 
to  form  it. 


80     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  law  of  definite  composition.  Attention  has  been 
called  to  the  fact  that  chemical  processes  involve  definite 
weights  of  matter  (p.  35).  We  have  now  seen  that  chem- 
ically pure  water  has  a  perfectly  definite  composition.  It 
makes  no  difference  what  the  source  of  the  water  is. 
Whether  it  is  obtained  from  the  melting  of  pure  ice,  from 
the  condensation  of  steam,  or  from  the  direct  combination 
of  hydrogen  and  oxygen,  its  composition  is  always  exactly 
the  same ;  namely,  1  part  by  weight  of  hydrogen  and 
7.94  parts  by  weight  of  oxygen,  or,  expressed  in  per- 
centages, 11.19  per  cent  of  hydrogen  and  88.81  per  cent 
of  oxygen.  A  similar  statement  has  been  found  to  hold 
good  for  every  pure  compound.  Thus,  experiments  have 
shown  that  potassium  chlorate  always  contains  31.90  per 
cent  of  potassium,  28.94  per  cent  of  chlorine,  and  39.16 
per  cent  of  oxygen,  while  mercuric  oxide  contains  92.61 
per  cent  of  mercury  and  7.39  per  cent  of  oxygen.  This 
truth  regarding  the  constancy  of  composition  of  chemical 
compounds  is  known  as  the  law  of  definite  composition, 
and  may  be  expressed  as  follows :  The  composition  of  a 
pure  compound  is  always  precisely  the  same. 

History  of  law  of  definite  composition.  The  common  expe- 
riences of  the  earlier  chemists  led  them  to  believe  that  the 
composition  of  a  pure  compound  is  quite  definite.  The  ques- 
tion as  to  whether  this  is  so  or  not  became  an  important  issue 
in  the  years  1802-1808,  as  a  result  of  the  views  of  a  distin- 
guished Frenchman,  Berthollet.  On  theoretical  grounds  Ber- 
thollet  was  led  to  believe  that  the  composition  of  a  substance 
is  somewhat  variable,  being  dependent  on  the  relative  quanti- 
ties of  the  several  materials  present  at  the  time  of  its  formation. 

These  views  were  strongly  opposed  by  a  fellow  countryman, 
Proust,  who  was  professor  of  chemistry  at  Madrid  during  most 
of  the  controversy.  Proust  maintained  that  the  composition 
of  a  pure  compound  is  perfectly  definite,  and  that  when  two 


WATER;  HYDROGEN  PEROXIDE  81 

elements  form  more  than  one  compound,  each  has  its  own  exact 
composition,  there  being  no  intermediate  gradations.  He  main- 
tained that  apparent  variability  is  due  to  lack  of  purity  in  the 
compound.  Proust's  experimental  work  was  very  accurate  for 
his  time,  but  his  analyses  were  subject  to  errors  of  from  1  to  2 
per  cent.  The  advance  in  experimental  exactness  has  steadily 
demonstrated  the  correctness  of  Proust's  conclusions.  In  1860 
and  again  in  1866  the  Belgian  chemist  Stas  undertook  elabo- 
rate researches  in  a  critical  study  of  the  law  of  definite  com- 
position, his  analyses  being  trustworthy  in  some  instances  to 
within  about  1  part  in  50,000.  Within  these  limits  he  showed 
that  the  law  holds  rigidly.  In  our  own  time  the  work  of  the 
American  chemist  Theodore  Richards  has  demonstrated  the 
accuracy  of  the  law  within  still  narrower  limits. 

Hydrogen  peroxide.  In  1818,  while  studying  the  action 
of  acids  upon  certain  oxides,  the  French  chemist  Thenard 
discovered  the  compound  which  we  now  call  hydrogen 
peroxide.  The  pure  compound  is  a  liquid  and,  like  water, 
is  composed  of  hydrogen  and  oxygen.  The  proportions 
in  which  the  hydrogen  and  oxygen  are  present  in  these 
two  compounds,  however,  are  widely  different,  as  shown 
in  the  following  statement: 

Water 1  part  of  hydrogen  to  7.94  parts 

of  oxygen  by  weight 

Hydrogen  peroxide 1  part  of  hydrogen  to  15.88  parts 

of  oxygen  by  weight 

In  other  words,  the  weight  of  oxygen  combined  with  a  fixed 
weight  of  hydrogen  is  just  twice  as  great  in  hydrogen  per- 
oxide as  in  water.  This  larger  percentage  of  oxygen  is 
indicated  by  the  name  peroxide,  the  prefix  per  meaning 
"  more  "  or  "  excess." 

Preparation  of  hydrogen  peroxide.  While  a  dilute  solu- 
tion of  hydrogen  peroxide  may  be  easily  obtained,  the  pure 
compound  cannot  be  prepared  without  great  difficulty, 


82     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

since  it  decomposes  into  water  and  oxygen  with  explosive 
violence.  Dilute  solutions  of  the  compound  are  prepared 
by  the  action  of  sulfuric  acid  on  barium  peroxide.  The 
reaction  may  be  represented  as  follows: 

sulfuric  acid  +  barium  peroxide — Hbarium  sulf ate  +  hydrogen  peroxide 
rhydrogen"!  rbarinm"!  rbarium"]  [hydrogen] 

sulfur  LoxygenJ  sulfur  [oxygen     J 

Loxygen    J  LoxygenJ 

It  will  be  noted  that  in  this  reaction  the  barium  of  the 
barium  peroxide  changes  places  with  the  hydrogen  of  the 
acid.  The  barium  sulfate  formed  is  insoluble,  while  the  hy- 
drogen peroxide  dissolves  in  the  water  present.  The  barium 
sulfate  may  therefore  be  removed  from  the  solution  by 
filtration.  In  this  way  one  can  readily  prepare  a  dilute 
solution  of  the  peroxide  in  water.  In  this  form  it  is  a 
common  article  of  commerce. 

Properties  of  hydrogen  peroxide.  Hydrogen  peroxide  is 
a  clear,  sirupy  liquid  having  a  density  of  1.458.  Because 
of  the  highly  explosive  character  of  the  pure  compound 
it  is  prepared  in  the  form  of  a  dilute  solution  in  water. 

Since  hydrogen  peroxide  so  readily  decomposes,  with 
evolution  of  oxygen,  it  acts  as  a  strong  oxidizing  agent, 
even  in  very  dilute  solutions.  An  easily  oxidizable  sub- 
stance, like  wool,  is  ignited  by  the  addition  of  a  few  drops 
of  the  pure  compound.  The  speed  of  decomposition  of 
hydrogen  peroxide  is  influenced  in  many  ways.  In  dilute 
solutions  and  at  a  low  temperature  the  speed  is  very  slow, 
while  at  higher  temperatures  and  in  more  concentrated 
solutions  it  becomes  so  great  as  to  cause  violent  explo- 
sions. Moreover,  the  speed  of  decomposition  is  greatly 
affected  by  the  presence  of  certain  catalytic  agents  (p.  26). 
Thus  a  little  finely  divided  platinum  or  manganese  dioxide, 
added  to  a  concentrated  solution  of  the  peroxide,  produces 


WATER;  HYDROGEN  PEROXIDE  83 

such  rapid  decomposition  as  to  cause  an  explosion.  Cer- 
tain organic  substances  have  a  similar  action.  Just  as  some 
substances  increase  the  rapidity  of  decomposition,  so  others 
retard  it.  Thus  the  ordinary  solution  of  hydrogen  peroxide 
sold  for  medicinal  purposes  contains  a  small  amount  of 
some  such  substance,  generally,  a  trace'  of  acid,  which  is 
added  to  preserve  the  strength  of  the  solution  by  retarding 
decomposition. 

Uses.  Hydrogen  peroxide  has  many  commercial  uses, 
all  based  on  its  strong  oxidizing  properties.  The  common 
medicinal  peroxide  of  the  druggist  is  an  aqueous  solution 
containing  3  per  cent,  by  weight,  of  the  compound.  It  has 
long  been  used  as  a  germicide,  but  recent  experiments  indi- 
cate that  its  efficiency  for  this  purpose  has  been  greatly 
overrated.  It  acts  upon  certain  dyes  and  natural  colors, 
such  as  that  of  the  hair,  oxidizing  them  into  colorless 
compounds ;  hence  it  is  sometimes  used  as  a  bleaching 
agent.  The  chemist  finds  it  especially  useful  as  an  oxi- 
dizing agent  in  many  analytical  operations.  For  this  pur- 
pose it  is  often  convenient  to  have  a  rather  concentrated 
solution,  so  that  a  30  per  cent  solution  is  now  sold  as 
a  commercial  product. 

The  law  of  multiple  proportion.  We  have  seen  that  both 
water  and  hydrogen  peroxide  are  compounds  of  hydrogen 
and  oxygen  and  that  the  ratio  by  weight  in  which  these 
two  elements  are  present  in  each  of  these  compounds  is 
as  follows: 

Water hydrogen  :  oxygen  =  1 :  7.04 

Hydrogen  peroxide  ....     hydrogen  :  oxygen  —  1 : 15.88 

It  will  be  seen  that  the  ratio  between  the  weights  of  oxy- 
gen combined  with  a  fixed  weight  of  hydrogen  (say  1  g.) 
in  these  two  compounds  is  7.94  : 15.88,  or  1 :  2. 


84     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Similarly,  many  elements  other  than  oxygen  and  hydro- 
gen unite  to  form  a  number  of  distinct  compounds,  each 
with  its  own  precise  composition.  In  all  such  compounds 
the  same  statement  holds  as  in  the  case  of  water  and  hydro- 
gen peroxide  —  the  weights  of  the  one  element  which  are 
combined  with  a  fixed  weight  of  the  other  always  bear  a 
simple  ratio  to  each  other,  such  as  1 :  2  or  2 :  3.  This  truth 
is  known  as  the  law  of  multiple  proportion.  It  was  formu- 
lated by  John  Dalton  (p.  89)  in  1808,  and  may  be  stated 
thus:  When  any  two  elements,  A  and  £,  combine  to  form 
more  than  one  compound,  the  weights  of  A  which  unite  with 
any  fixed  weight  of  B  bear  the  ratio  of  small  whole  numbers 
to  each  other. 

EXERCISES 

1.  Why  does  the  chemist  use  distilled  water,  rather  than  filtered 
water,  in  making  solutions  ? 

2.  How  could  you  determine  the  total  amount  of  solid  matter 
dissolved  in  a  sample  of  water  ? 

3.  How  could  you  determine  whether  a  given  sample  of  water 
is  distilled  water? 

4.  How  could  the  presence  of  air  dissolved  in  water  be  detected  ? 

5.  How  could  the  amount  of  water  in  a  food  such  as  'bread  or 
potato  be  determined? 

6.  Would  ice  frozen  from  impure  water  necessarily  be  free  from 
disease  germs? 

7.  Suppose  that  the  maximum  density  of  water  were  at  0°  in- 
stead of  4° ;  what  effect  would  this  have  on  the  formation  of  ice  on 
bodies  of  water? 

8.  In  the  experiment  illustrated  by  Fig.  31,  why  is  cold  water 
passed  into  C  instead  of  into  D? 

9.  Mention  at  least  two  advantages  that  a  metal  condenser  has 
over  a  glass  condenser. 

10.  Draw  a  diagram  of  the  apparatus  used  in  your  laboratory  for 
supplying  distilled  water. 


WATER;  HYDROGEN  PEROXIDE  85 

11.  Report  upon  the  method  of  purifying  the  water  supply  in 
your  home  town  or  city. 

12.  Determine  the  ratio  in  which  oxygen  and  hydrogen  unite 
from  the  following  data,  the  volumes  all  being  measured  under  the 
same  conditions  of  temperature  and  pressure  : 

Volume  of  oxygen  in  eudiometer 8.54  cc. 

Volume  of  oxygen  and  hydrogen 52.72  cc. 

Volume  of  gas  (hydrogen)  left  after  explosion       .     .     27.10cc. 

13.  Morley  found  the  composition  of  water  by  determining  the 
weights  of  hydrogen  and  of  oxygen  that  combine  with  each  other 
to  form  water.    The  results  of  six  trials  are  as  follows : 
HYDROGEN  USED        OXYGEN  USED  HYDROGEN  USED      OXYGEN  USED 

(1)  3.2645 g.  25.9176  g.  (4)  3.8450 g.  30.5294  g. 

(2)  3.2559  g.  25.8531  g.  (5)  3.8382  g.  30.4700  g. 

(3)  3.8193  g.  30.3210  g.  (6)  3.8523  g.  30.5818  g. 

In  each  case  calculate  the  ratio  in  which  the  hydrogen  and  oxygen 
combined  to  form  water. 

14.  20  cc.  of  hydrogen  and  7  cc.  of  oxygen  are  placed  in  a  eudi- 
ometer and  the  mixture  exploded  as  in  Fig.  37.    (a)  How  many 
cubic  centimeters  of  water  vapor  are  formed?   (b)  What  gas  and 
how  much  of  it  remains  in  excess? 

15.  (a)  What  weight  of  water  can  be  formed  by  the  combustion  of 
100  1.  of  hydrogen,  measured  under  standard  conditions  ?    (6)  What 
volume  of  oxygen  would  be  required  in  (a)  ?    (c)  What  weight  of 
potassium  chlorate  is  necessary  to  prepare  this  amount  of  oxygen  ? 

16.  What  weight  of  oxygen  is  present  in  1  kg.  of  the  ordinary 
hydrogen  dioxide  solution  ?   In  the  decomposition  of  this  weight  of 
the  dioxide  into  water  and  oxygen,  what  volume  of  oxygen  (meas- 
ured under  standard  conditions)  is  evolved  ? 


CHAPTER  VII 
COMBINING  WEIGHTS;   THE  ATOMIC  THEORY 

Introduction.  We  have  already  considered  three  laws 
which  deal  with  the  relations  by  weight  which  hold  good 
during  chemical  action :  (1)  the  law  of  conservation  of 
matter,  (2)  the  law  of  definite  composition,  (3)  the  law 
of  multiple  proportion.  To  these  must  now  be  added  a 
fourth,  —  the  law  of  combining  weights. 

Combining  weights.  We  have  seen  that  hydrogen  and 
oxygen  combine  in  two  perfectly  definite  ratios  by  weight ; 
namely,  1  g. :  7.94  g.  and  1  g. :  (2  x  7.94  g.).  In  a  similar 
way  it  is  easy  to  determine  the  ratios  in  which  elements 
other  than  oxygen  combine  with  hydrogen.  For  example, 
hydrogen  combines  with  sulfur  to  form  a  gas  called  hy- 
drogen sulfide,  and  with  the  metal  calcium  to  form  a  solid 
called  calcium  hydride.  In  these  compounds  the  ratios  by 
weight  are 

Water hydrogen  1  g.,  oxygen  7.94  g. 

Hydrogen  peroxide   .     .  hydrogen  1  g.,  oxygen  2  x  7.94  g. 

Hydrogen  sulfide  .     .     .  hydrogen  1  g.,  sulfur  16  g. 

Calcium  hydride  .     .     .  hydrogen  1  g.,  calcium  19.88  g. 

It  is  therefore  possible  to  assign  to  each  element  combin- 
ing with  hydrogen  a  number  which  expresses  the  weight 
in  grams  of  the  element  which  combines  with  1  g.  of 
hydrogen. 

Now  experiment  reveals  a  very  interesting  fact.  The 
numbers  which  express  the  ratios  in  which  two  elements 


COMBINING  WEIGHTS;  THE  ATOMIC  THEORY  87 

combine  with  a  fixed  weight  of  hydrogen  also  express  the 
ratio  in  which  they  combine  with  each  other.    Thus, 

2  x  7.94  g.  of  oxygen  combines  with  16  g.  of  sulfur 
7.94  g.  of  oxygen  combines  with  19.88  g.  of  calcium 
19.88  g.  of  calcium  combines  with  16  g.  of  sulfur 

From  a  study  of  a  great  number  of  combining  ratios 
like  those  just  given  it  has  been  determined  that  elements 
do  not  combine  in  simple  ratios  by  weight,  say  1  g.  of 
one  element  with  1  g.  or  with  2  g.  of  another ;  but  each 
element  has  a  weight  peculiar  to  itself  that  expresses  how 
much  of  that  element  will  combine  with  any  other  ele- 
ment. These  numbers  are  called  the  combining  weights 
of  the  elements. 

Standard  for  combining  weights.  The  numbers  assigned 
to  oxygen,  calcium,  and  sulfur  in  the  last  paragraph  were 
determined  by  finding  how  many  grams  of  each  of  these 
elements  combine  with  1  g.  of  hydrogen.  They  are  there- 
fore relative  to  hydrogen  taken  as  unity.  Any  other  ele- 
ment might  have  been  taken  as  the  standard,  and  instead  of 
1  g.  of  the  standard  element  any  convenient  weight  might 
have  been  taken.  The  weights  found  are  ratio  weights, 
and  while  the  weights  will  all  be  changed  if  we  change 
the  standard,  their  ratios  to  each  other  will  remain  the  same. 
For  many  reasons  it  is  better  to  select  oxygen  rather  than 
hydrogen  as  the  standard  element  It  is  likewise  better  to 
select  8  g.  rather  than  1  g.  as  its  standard  value,  so  that 
no  other  element  may  have  a  combining  weight  of  less 
than  unity.  If  oxygen  is  taken  as  8,  hydrogen  becomes 
1.008 ;  calcium,  20.03 ;  and  sulfur,  16.03. 

The  combining  weight  of  an  element  may  therefore  be 
defined  as  that  weight  of  the  element  which  will  combine 
with  8  g.  of  oxygen. 


88     AN  ELEMENTAKY  STUDY  OF  CHEMISTRY 


In  Fig.  38  the  lines  connecting  any  two  symbols  at  once 
indicate  the  ratio  by  weight  in  which  the  elements  combine. 
Elements  with  more  than  one  combining  weight.  It  is 
evident  that  some  elements  have  more  than  one  combining 
weight,  for  we  have  seen  that  1  part  of  hydrogen  combines 
either  with  7.94  or  with  15.88  parts  of  oxygen.  Or,  if  we 
take  8  g.  of  oxygen  as  standard,  the  combining  weight  of 
hydrogen  may  be  either  1.008  g.  or  0.504  g.  In  all  suck  cases 
the  number  expressing  the  larger  combining  weight  is  a  simple 
multiple  of  the  number  expressing  the  smaller  combining  weight. 
The  law  of  combining  weights. 
1.008  The  law  of  combining  weights 
may  now  be  stated  as  follows: 
To  each  element  may  be  assigned 
a  number  which  in  itself,  or  when 
multiplied  by  some  integer,  expresses 
the  weight  by  which  the  element 
combines  with  other  elements. 

The  atomic  theory.  The  four 
laws  which  in  general  terms  state 
the  way  in  which  elements  combine 
with  each  other  are  remarkably 

simple.  We  instinctively  feel  that  the  reason  for  this  sim- 
plicity must  lie  in  the  way  matter  is  made  up,  and  we 
begin  to  imagine  a  structure  for  matter  that  if  it  were  real 
would  result  in  these  simple  laws.  In  other  words,  we  try 
to  devise  a  theory  of  the  structure  of  matter  and  the 
mechanism  of  chemical  action.  Of  all  the  theories  that 
have  been  advanced  concerning  the  nature  of  matter,  the 
one  proposed  by  John  Dalton  (Fig.  39),  and  known  as  the 
atomic  tlieory,  is  the  most  satisfactory.  The  main  points 
of  this  theory  in  its  present  form,  together  with  the  reasons 
for  making  them  a  part  of  the  theory,  are  as  follows : 


FIG.  38.    Diagram   showing 
the  combining  ratios  of  oxy- 
gen, hydrogen,  sulfur,  and 
calcium 


COMBINING  WEIGHTS;  THE  ATOMIC  THEORY  89 

1.  Every   weighable    quantity   of    an    elementary   sub- 
stance is  made  up  of  a  very  great  number  of  unit  bodies 
called  atoms. 

2.  Experiment  shows  that  the  composition  of  a  given 
compound  is  always  the  same  (law  of  definite  composi- 
tion).    The  simplest  way  to  adjust  the  theory  of  atoms 
to  this  fact  is  to  assume  that  the  atoms  of  each  element 
all   have  the  same  weight, 

while  those  of  different  ele- 
ments have  different  weights, 
and  that  during  chemical 
union  a  definite  number  of 
one  kind  of  atoms  combines 
with  a  definite  number  of 
another  kind  to  form  a 
particle  of  a  compound.  If 
this  should  be  true,  a  given 
compound  would  of  neces- 
sity have  a  perfectly  definite 
composition. 

3.  Since  there  is  no  change 

in    weight  when  two    sub-     FlG" 39'  John  Dalton  (1766-1844) 
stances  act  upon  each  other,     The  E,nslish  school-teacher  and  chem- 

ist  who  suggested  the  atomic  theory 

the  weights  of  the  individual 

atoms  must  remain  unchanged  as  a  result  of  the  action. 

4.  To   account  for  the  law  of  multiple  proportion  we 
must  assume  that  the   atoms  of  two  different   elements 
may  unite  in  different  ratios ;    for  example,  if  one  atom 
of  A  unites  with  one  of  B  under  one  set  of  conditions, 
but  with  two  of  B  under  other  conditions,  then  we  shall 
have  two  different  compounds.     The  masses  of  B  com- 
bined with  a  fixed  mass  of  A  will  be  in  the  ratio  of  1:2, 
since  the  number  of  atoms  are  in  this  ratio. 


90     AN  ELEMENTARY  STUDY  OF  CHEMISTEY 

5.  The  law  of  combining  weights  tells  us  that  a  defi- 
nite number  can  be  assigned  to  each  element,  which  ex- 
presses its  combining  value.  If  each  atom  has  its  own 
peculiar  weight  and  if  atoms  always  combine  with  each 
other  in  definite  numbers,  then  these  combining  numbers 
indicate  the  relative  weights  of  the  atoms  themselves.  That 
an  element  may  have  two  different  combining  weights, 
one  a  multiple  of  the  other,  is  provided  for  by  the  sup- 
position that  the  atoms  are  able  to  combine  in  several 
different  ratios. 

Summary  of  the  atomic  theory.  The  atomic  theory  sug- 
gests that  all  matter  is  made  up  of  minute  bodies  called 
atoms.  The  atoms  of  a  given  element  are  all  alike  in 
weight,  but  those  of  different  elements  have  different 
weights.  When  elements  act  upon  each  other,  the  action 
takes  place  between  definite  small  numbers  of  the  atoms. 

While  the  atomic  conception  of  matter  is  still  referred 
to  as  the  atomic  theory,  so  many  facts  are  now  known 
that  are  in  complete  harmony  with  this  view  that  it  can 
hardly  be  doubted  that  this  theory  gives  us  a  picture  of 
real  facts  and  that  the  reality  of  atoms  and  molecules  is 
virtually  a  matter  of  conclusive  proof. 

Molecules  and  atoms.  Dalton  applied  the  name  atom  to 
both  elements  and  compounds.  It  is  evident,  however, 
that  the  smallest  particle  of  a  compound  must  be  made 
up  of  at  least  two  different  kinds  of  atoms.  Thus  the 
smallest  particle  of  water  that  can  be  formed  must  con- 
tain both  hydrogen  and  oxygen ;  such  a  particle  of  water 
is  called  a  molecule  of  water,  and  any  amount  of  water  is 
simply  a  collection  of  a  definite  number  of  these  molecules. 
In  general,  it  may  be  said  that  the  term  molecule  is  now 
applied  to  the  smallest  particle  which,  taken  in  large  num- 
bers, makes  up  the  bodies  with  which  we  deal. 


COMBINING  WEIGHTS;  THE  ATOMIC  THEORY  91 

Molecules  of  elements.  Since  two  kinds  of  atoms  unite 
to  form  a  molecule  of  a  compound,  the  question  naturally 
arises,  Do  two  or  more  atoms  of  the  same  kind  combine 
to  form  a  molecule  of  an  elementary  substance,  or  do  ele- 
mentary substances  consist  of  separate  atoms  ?  It  has 
been  found  that  the  elements  differ  among  themselves  in 
this  respect.  In  some  cases,  as  with  mercury,  the  atoms 
do  not  unite,  so  that  the  molecule  is  the  same  as  the 
atom;  in  other  cases,  as  with  oxygen  and  hydrogen,  two 
atoms  unite  to  form  a  molecule  of  the  element.  The  ex- 
periments which  prove  that  these  statements  are  true  will 
be  described  later. 

Atomic  weights.  It  would  be  of  great  interest  if  we 
could  determine  the  actual  weights  of  the  various  kinds 
of  atoms.  They  are  so  very  small,  however,  that  we  can 
never  hope  to  determine  their  weight  even  approximately. 
It  has  been  shown  that  the  smallest  particle  visible  with 
the  most  powerful  microscope  ever  constructed  contains 
at  least  1000  atoms. 

We  have  seen,  however,  that  the  ratio  between  the 
combining  weights  is  the  same  as  that  between  the  weights 
of  the  atoms  themselves,  so  we  should  be  able  to  deter- 
mine their  relative  weights  with  precision.  But  most  of 
the  elements  have  more  than  one  combining  weight,  and 
we  must  find  some  means  of  choosing  the  one  which 
correctly  expresses  the  weight  of  a  single  atom. 

Before  this  problem  can  be  solved,  methods  must  be 
devised  for  finding  the  relative  weights  of  molecules  of  com- 
pounds. Such  methods  have  been  developed  and  will  be 
described  later  (Chap.  XXIII).  These  methods  have  led 
to  the  adoption  of  a  single  number  for  each  element,  called 
its  atomic  weight.  A  list  of  atomic  weights  will  be  found 
on  the  back  cover  of  the  book. 


92     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

EXERCISES 

1.  Report  the  important  events  in  the  life  of  Dal  ton  (consult 
encyclopedia). 

2.  State  the  four  laws  mentioned  in  this  chapter. 

3.  Calculate  the  weights  of  hydrogen  that  combine  with  100  g. 
of  oxygen  to  form  water  and  hydrogen  peroxide  respectively.    Are 
your  results  in  accord  with  the  law  of  multiple  proportion  ? 

4.  What  is  the  derivation  and  meaning  of  the  word  atom  ?   (Con- 
sult dictionary.) 

5.  Why  did  Dalton  make  the  assumption  that  the  atoms  of  the 
same  element  all  have  exactly  the  same  weight  ? 

6.  What  is  meant  by  the  following  terms :  (a)  atom  of  an  ele- 
ment ;  (&)  molecule  of  an  element ;   (c)  molecule  of  a  compound  ? 

7.  Would  it  be  logical  to  speak  of  an  atom  of  a  compound? 


CHAPTER  VIII 
FORMULAS;   EQUATIONS;   CALCULATIONS 

Percentage  composition.  Just  as  we  can  determine  the 
composition  of  water  with  great  accuracy,  so,  by  similar 
means,  we  can  determine  the  composition  of  other  com- 
pounds. Having  analyzed  a  given  compound,  we  usually 
express  its  composition  hi  percentages,  or  in  the  parts  of 
each  element  present  in  100  parts  of  the  compound.  Thus, 
we  have  seen  that  water  consists  of  88.81  per  cent  of 
oxygen  and  11.19  per  cent  of  hydrogen. 

Formulas.  The  plan  of  stating  composition  in  percent- 
ages takes  no  account  of  the  fact  that  compounds  are 
made  up  of  molecules  in  which  each  atom  has  a  charac- 
teristic weight.  It  would  be  much  better  to  have  a  method 
of  stating  composition  which  would  express  all  these  facts. 

Since  the  molecule  of  any  chemical  compound  consists  of 
a  definite  number  of  atoms,  it  is  very  convenient  to  repre- 
sent the  composition  of  a  compound  by  indicating  the  num- 
ber and  the  variety  of  atoms  that  make  up  its  molecules. 
This  can  be  done  by  the  use  of  symbols  of  the  elements.  In 
this  way  HgO  will  represent  mercuric  oxide,  a  molecule  of 
which  has  been  found  to  contain  one  atom  of  mercury  (Hg) 
and  one  of  oxygen  (O).  Similarly,  H2O  will  represent 
water,  the  molecules  of  which  consist  of  two  atoms  of  hydro- 
gen and  one  of  oxygen,  the  subscript  figure  indicating  the 
number  of  atoms  of  the  element  whose  symbol  it  follows. 
These  combinations  of  symbols,  which  represent  the  atomic 
composition  of  molecules  of  substances,  are  called  formulas. 


94     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Formulas  from  percentage  composition.  If  we  know  the 
percentage  composition  of  a  compound  as  the  result  of 
analysis  and  can  devise  a  method  for  determining  the 
molecular  weight  of  the  compound,  it  will  be  easy  to 
calculate  its  formula.  Thus,  if  we  know  that  potassium 
chlorate  contains  31.90  per  cent  of  potassium  (K),  28.94  per 
cent  of  chlorine  (Cl),  and  39.16  per  cent  of  oxygen  (O), 
and  that  taking  the  weight  of  the  oxygen  atom  as  16,  the 
weight  of  a  molecule  of  potassium  chlorate  is  122.56,  we 
may  reason  as  follows: 

The  portion  of  potassium  in  each  molecule  is  122.56 
X  31.90  per  cent  =  39.10  K ;  the  portion  of  chlorine  is 
122.56  X  28.94  per  cent  =  35.46  Cl;  the  portion  of  oxygen 
is  122.56  x  39.16  per  cent  =  48.00  O.  Now  the  weight  of 
the  molecule  (122.56)  is  the  sum  of  the  weights  of  the 
atoms  that  compose  it.  Therefore  the  numbers  we  have 
obtained  are  the  weights  of  the  several  kinds  of  atoms 
composing  the  molecule,  and  these  must  be  either  the 
weights  of  the  single  atoms  or  of  multiples  of  them. 
Reference  to  the  table  of  atomic  weights  shows  that  the 
numbers  obtained  for  potassium  and  for  chlorine  corre- 
spond to  the  weights  for  single  atoms ;  while  the  weight 
found  for  oxygen  is  that  of  three  atoms.  Consequently 
the  formula  for  potassium  chlorate  is  KClOg.  Later  we 
shall  find  general  methods  for  obtaining  molecular  weights 
of  compounds,  and  the  formulas  of  the  compounds  can  then 
be  calculated  by  the  above  method. 

Calculation  of  formulas  when  the  molecular  weight  is  not 
known.  If  we  do  not  know  the  molecular  weight  of  the  com- 
pound, but  merely  have  the  percentage  composition  as  found 
by  analysis,  we  can  determine  the  simplest  formula  the  com- 
pound can  have,  but  we  cannot  be  sure  that  this  is  the  true 
formula.  As  an  example,  let  us  suppose  that  we  have  found 


FORMULAS;  EQUATIONS;  CALCULATIONS     95 

the  percentage  composition  of  potassium  chlorate,  but  that  we 
do  not  know  that  its  molecular  weight  is  122.56. 

The  percentage  numbers  31.90  per  cent  K,  28.94  per  cent  01, 
and  39.16  per  cent  O  are  the  numbers  of  grams  of  these  ele- 
ments present  in  a  purely  arbitrary  weight  of  potassium  chlo- 
rate ;  namely,  in  100  g.  But  we  have  seen  that  a  gram  of  one 
element  does  not  as  a  rule  combine  with  a  gram  of  another 
element,  for  they  combine  in  the  ratio  of  their  atomic  weights. 
Consequently  the  amount  of  each  element  present  in  100  g.  of 
a  compound  bears  no  simple  relation  to  the  atomic  weights  of 
the  elements,  and  we  wish  to  change  the  percentage  numbers 
into  others  that  will  express  the  ratio  between  the  number  of 
the  several  kinds  of  atoms  present  in  100  g.  If  we  divide  the 
percentage  number  of  each  element  by  its  atomic  weight,  the 
quotients  will  be  the  relative  number  of  each  kind  of  atom 
present  in  100  g.  of  the  compound. 

31.90  H-  39.10  =  0.8158  =  relative  number  of  atoms  of  K  in  100  g. 
28.94  -^  35.46  =  0.8161  =  relative  number  of  atoms  of  Cl  in  100  g. 
39.16  -7-16  =  2.4475  =  relative  number  of  atoms  of  O  in  100  g. 

Now  100  g.  of  the  compound  is  made  up  of  a  great  number  of 
molecules,  and  the  ratio  of  the  various  atoms  in  a  single  mole- 
cule must  be  the  same  as  in  any  weight  of  the  compound.  But 
in  the  molecule  there  cannot  be  fractions  of  atoms,  but  only 
the  whole  numbers.  So  we  reduce  the  ratio  0.8158 :  0.8161 : 
2.4475  to  a  ratio  of  whole  numbers  by  dividing  the  three  num- 
bers by  the  least  of  the  three.  This  gives  us  the  ratio  1 : 1  :'3. 
If  this  is  the  correct  ratio  of  the  number  of  the  several 
kinds  of  atoms  in  the  molecule,  the  simplest  formula  is  KC10g. 
Evidently  the  formula  might  be  any  multiple  of  this  simplest 
formula  such  as  K2C1206  or  K3C1809,  for  compounds  represented 
by  any  of  these  formulas  would  all  give  the  same  percentage 
composition.  We  should  have  to  know  the  weight  of  the  mole- 
cule to  decide  between  these  possibilities.  Later  we  shall  find 
ways  to  show  that  the  molecular  weight  of  potassium  chlo- 
rate is  somewhere  near  120,  so  that  the  true  formula  is  the 
simplest  one ;  namely,  KC108  (122.56). 


96     AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Since  the  results  of  analysis  are  always  subject  to  small 
errors,  the  percentage  composition  of  a  compound  is  never 
known  from  experiment  with  entire  precision.  Consequently 
the  ratio  for  the  number  of  atoms  never  comes  out  exactly  an 
integer  ratio,  but  it  always  comes  so  near  to  one  as  to  leave  no 
doubt  as  to  the  real  ratio. 

Facts  expressed  by  formulas.  From  what  has  been  said, 
it  will  be  seen  that  formulas  are  used  to  express  several 
distinct  facts : 

1.  Atomic  composition  of  molecules.    A  formula  shows  the 
number  and  kinds  of  atoms  in  a  molecule  of  a  compound. 
The  formula  H2O  states  that  a  molecule  of  water  is  com- 
posed of  two  atoms  of  hydrogen  and  one  of  oxygen.    The 
formula  of  hydrogen  sulfate  (H2SO4)  shows  that  its  mole- 
cule consists  of  two  atoms  of  hydrogen,  one  of  sulfur,  and 
four  of  oxygen. 

2.  Molecular  weights  of  compounds.    Since  each  atom  has 
its  own  weight,  the  sum  of  all  the  atoms  in  a  molecule 
must  be  the  weight  of  the  molecule  itself  relative  to  oxy- 
gen taken  as  16.    The  relative  weight  of  the  molecule  of 
water  is  therefore  (2  x  1.008)  +  16  =  18.016.   The  relative 
weight  of  the  molecule  of  hydrogen  sulfate   (H2SO4)  is 
(2  x  1.008)  +  32.06  +  (4  x  16)  =  98.076. 

•3.  Percentage  composition  of  compounds.  From  a  formula 
we  can  easily  go  back  to  the  percentages  from  which  it  was 
calculated.  Thus,  if  the  molecule  of  water  weighs  18.016 
and  contains  one  oxygen  atom  of  weight  16,  the  frac- 

1  f\ 

tion  of  its  weight  due  to  oxygen  is  —     — ,  or  88.81  per  cent. 

18.01b 

The  fraction  due  to  hydrogen  is      '          =11.19  per  cent. 

18.016 

Gram-molecular  weights ;  formula  weights.  For  practi- 
cal purposes  we  deal  with  pounds  or  with  grams  of  a 


FORMULAS;  EQUATIONS;  CALCULATIONS     97 

substance,  not  with  atoms  and  molecules.  Now,  since  the 
numbers  18.016,  16,  and  2.016  represent  the  ratio  by 
weight  between  a  molecule  of  water  and  the  oxygen  and 
hydrogen  of  which  it  is  composed,  the  same  ratios  must 
hold  between  any  weight  of  water  we  may  choose  and 
the  oxygen  and  hydrogen  in  this  weight  of  water,  for 
any  amount  of  water  is  simply  a  collection  of  molecules 
of  water.  Evidently,  in  18.016  Ib.  of  water  there  will 
be  16  Ib.  of  oxygen  and  2.016  Ib.  of  hydrogen,  and  in 
18.016  g.  there  will  be  16  g.  of  oxygen  and  2.016  g.  of 
hydrogen. 

For  practical  purposes,  therefore,  we  may  allow  the 
symbol  H  to  stand  for  1.008  grams  of  hydrogen,  the  sym- 
bol O  for  \Qgrams  of  oxygen,  and  the  formula  H2O  for 
18.016  grams  of  water.  The  weight  in  grams  of  an  ele- 
ment, corresponding  to  its  atomic  weight,  is  called  a 
gram-atomic  weight  or  symbol  weight.  The  weight  in  grams 
of  a  compound,  corresponding  to  its  molecular  weight,  is 
called  a  gram-molecular  weight  or  formula  weight. 

Equations.  Having  devised  a  convenient  way  of  ex- 
pressing the  composition  of  compounds,  not  in  percent- 
ages but  in  formulas,  we  make  use  of  equations  to  express 
chemical  transformations,  using  an  arrow  in  place  of  an 
equality  sign.  For  example,  the  equation 

2H+0 )-H2O  (1) 

is  a  concise  method  of  stating  two  distinct  facts. 

1.  Qualitatively,  it  states  that  water  is  formed  by  the 
union  of  hydrogen  and  oxygen. 

2.  Quantitatively,  it  tells  us  that  2  symbol  weights  of 
hydrogen    (2.016  g.)    combine  with   1   symbol  weight  of 
oxygen    (16  g.)    to    form    a    formula    weight    of    water 
(18.016  g.). 


98     AN  ELEMENTAKY  STUDY  OF  CHEMISTRY 

Molecular  equations.  Since  a  formula  expresses  the 
composition  of  a  molecule,  and  since  experiment  kas 
shown  that  a  molecule  of  oxygen  and  one  of  hydrogen 
each  contain  two  atoms,  the  formulas  of  these  gases  are 
written  O2  and  H2  rather  than  2  O  or  2  H,  which  would 
simply  represent  two  atoms  not  combined.  If  we  wish 
equation  (1)  to  state  these  additional  facts,  we  shall 
have  to  change  it  to  the  form 

2II2  +  02 — J-2H.O  (2) 

This  is  called  a  molecular  equation,  and  it  will  be  seen 
that  it  expresses  the  same  ratios  by  weight  as  does  equa- 
tion (1).  It  also  expresses  the  fact  that  2  molecules  of 
hydrogen  combine  with  1  molecule  of  oxygen  to  form 
2  molecules  of  water,  and  this  makes  it  a  more  useful 
equation. 

Decomposition  of  potassium  chlorate.  Let  us  take  another  ex- 
ample. It  will  be  remembered  that  oxygen  was  prepared  by 
heating  potassium  chlorate,  which  has  the  formula  KC10g. 
When  heated,  this  compound  decomposes  into  oxygen  and  a 
compound  called  potassium  chloride,  whose  formula  is  KC1. 
The  decomposition  is  represented  by  the  equation 

2  KC108 >-  2  KC1  +  3  02 

This  equation  states  the  following  facts  : 

1.  Qualitatively,  potassium  chlorate  decomposes  into  potas- 
sium choride  and  oxygen. 

2.  Quantitatively,  2  formula  weights  of  potassium  chlorate 
(2  x  122.6  g.)  decompose   into   2  formula  weights  of  potas- 
sium chloride  (2  x  74.69  g.)  and  3  formula  weights  of  oxygen 
(3  x  32  g.).    The  coefficient  before  a  formula  applies  to  the 
formula  as  a  whole,  while  the  subscript  number  applies  only 
to  the  symbol  which  it  follows. 

3.  Molecularly,  2  molecules  of  potassium  chlorate  decompose 
into  2  molecules  of  potassium  chloride  and  3  of  oxygen. 


FORMULAS;  EQUATIONS;  CALCULATIONS     99 

Equations  of  reactions  so  far  studied.  Let  us  now  put 
into  the  form  of  equations  a  number  of  the  reactions 
studied  up  to  this  point,  remembering  that  all  of  these 
equations  rest  upon  careful  experimental  analysis. 

1.  Preparation  of  oxygen: 

From  mercuric  oxide: 
2HgO 
From  potassium  chlorate  : 

2KC1O3 
From  the  electrolysis  of  water: 

2H20  —  ^2H2  +  02 
From  sodium  peroxide  and  water  : 

2  Na2O2  +  2  H2O  —  *-  4  NaOH  +  Ot 

2.  Preparation  of  hydrogen  : 
From  sodium  and  water: 


From  zinc  and  sulfuric  acid  : 

Zn  +  H2SO4  -  >-  ZnS04  +  H2 
From  iron  and  sulfuric  acid: 

Fe  -f  H2SO4  —  *•  FeSO4  4-  H2 
From  steam  and  iron: 

3  Fe  +  4  H2O  —  »•  Fe8O4  +  4  H2 

3.  Preparation  of  hydrogen  peroxide  : 

BaO2  +  H2SO4  —  *•  BaSO4  +  H2O2 

Representation  of  the  heat  of  reaction.  We  can  also 
employ  chemical  equations  to  express  the  heat  given  off 
or  absorbed  during  chemical  action.  The  equation 

*  2  H2O  +  138,000  cal. 


100    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

states  the  fact  that  when  4.032  g.  of  hydrogen  combines 
with  32  g.  of  oxygen,  forming  36.032  g.  of  water,  heat 
is  given  off  to  the  extent  of  138,000  cal.  Evidently, 
when  1  formula  weight  (18.016  g.)  of  water  is  formed, 
69,000  cal.  is  given  off,  and  this  is  called  the  heat  of 
formation  of  water. 

Conditions  of  a  reaction  not  indicated  by  equations. 
Equations  merely  state  the  composition  of  the  substances 
taking  part  in  the  reaction  and  the  weights  of  each  one 
involved,  together  with  the  energy  change  measured  as 
heat.  They  do  not  tell  the  conditions  under  which  the 
reaction  will  take  place.  For  example,  the  equation 

2HgO  —  ^2Hg  +  02 

does  not  tell  us  that  it  is  necessary  to  keep  heating  the 
mercuric  oxide  to  a  moderately  high  temperature  in  order 
to  effect  its  decomposition.  The  equation 

Zn  +  H2SO4  -  +•  ZnSO4  +  H2 

in  no  way  indicates  that  the  hydrogen  sulfate  must  be  dis- 
solved in  water  before  it  will  act  upon  zinc.  The  equation 


does  not  indicate  that  no  perceptible  action  takes  place  un- 
less the  sulfur  is  first  heated,  but  that  when  once  started 
it  goes  on  of  its  own  accord  and  with  a  bright  flame. 

It  will  therefore  be  necessary  to  pay  close  attention  to 
the  details  of  the  conditions  under  which  a  given  reaction 
occurs,  as  well  as  to  the  statement  of  the  equation  itself. 

Problems  based  on  equations.  Since  an  equation  is  a 
statement  of  the  weights  of  materials  which  take  part  in 
a  reaction,  when  the  equation  has  once  been  established 
by  experiment  we  can  use  it  in  calculating  the  various 
weights.  A  few  examples  will  show  how  this  may  be  done. 


FORMULAS;  EQUATIONS;  CALCULATIONS   101 

1.  How  many  grains  of  oxygen  will  be  evolved  on  heating 
100  g.  of  mercuric  oxide  ? 

First  write  the  equation  for  the  reaction  involved: 

2HgO ^2Hg  +  02  (1) 

Next  determine  the  relative  weights  of  the  amounts  of  the 
different  substances  involved  in  the  reaction.  The  atomic 
weights  of  mercury  and  oxygen  are  respectively  200.6  and  1(5 
(see  table  on  back  cover).  Hence  the  relative  weight  of  the 
2HgO  equals  2(200.6  +  16),  or  433.2.  Similarly,  the  relative 
weight  of  the  oxygen  evolved,  namely,  02,  equals  2  x  16,  or 
32.  It  is  convenient  now  to  write  these  numbers  under  the 
formulas  in  equation  (1).  This  then  becomes 


2  HgO  >•   -   J.i&    -p-    x-r- 

433.2  32 

These  numbers  indicate  that  433.2  units  by  weight  (in  this  case 
grams)  of  mercuric  oxide  will,  on  being  heated,  evolve  32  units 
by  weight  of  oxygen ;  hence  1  g.  of  mercuric  oxide  will  give 

32  32 

433-2  &•  of  oxygen,  and  100  g.  will  give  100  x  jggTj,  or  7.38  g.; 

or  the  relation  between  the  weights  of  the  substances  involved 
may  be  stated  in  the  form  of  a  proportion  : 

433.2      100 

32    :=    x' 
x  =  7.38 

2.  I  wish  to  prepare  100 g.  of  oxygen,  using  potassium  chlorate 
as  a  source  of  the  oxygen.    How  many  grams  of  the  chlorate 
will  be  required  ? 

2  KC108 *•  2  KC1  +  30. 

245.12  96 

Proportion  :         '      =  — — ;  or  x  =  255.33  g. 

3.  How  many  grams  of  zinc  must  be  dissolved  in  sulfuric 
acid  to  produce  10  g.  of  hydrogen  ? 

Zn  +  H,S04 >•  ZnSO.  +  H2 

65.37  .  2.016 


102    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Proportion  :  ^jg  =  ^  ;  or  x  =  324.2  g. 

It  must  be  remembered  that  the  equations  show  relations 
by  weight,  not  by  volume ;  hence  in  problems  involving  vol- 
umes of  gases  it  will  be  necessary  first  to  find  the  weights  of 
the  gases.  The  table  in  the  Appendix  gives  the  weight  of  1  1. 
of  each  of  the  common  gases,  measured  under  standard  con- 
ditions. The  following  problem  will  illustrate  the  method : 

4.  How  many  grams  of  potassium  chlorate  are  necessary  to 
prepare  100  1.  of  oxygen  ? 

Since  1 1.  of  oxygen  weighs  1.429  g.,  100 1.  will  weigh  142.9  g. 

2KC103 >-2KCl-f-302 

245.12  96 

Proportion  :  ^-  =  .mx  =  363.8  g. 


EXERCISES 

1.  State  all  the  facts  expressed  by  the  formulas  HC1,  IINO3, 
Ca(OH)2,  and  H3PO4. 

.2.  Calculate  the  percentage  composition  of  the  following  com- 
pounds from  the  formulas  as  given  :  (a)  potassium  chlorate  (KC1O3); 
(&)  hydrogen  sulfate  (H2SO4);  (c)  water  (H2O);  (d)  saltpeter  (KNO3); 
(e)  baking  soda  (NaHCO,). 

3.  From  t%he  following  analyses  calculate  the  simplest  formula  : 

(1)  8  =  39.07%  0  =  58.49%  H=    2.44% 

(2)  Ca  =  29.40%  S  =  23.56%  O  =  47.04% 
(8)     K  =  38.67%               N  =  13.88%  O  =  47.45% 

4.  It  is  required  to  prepare  30  g.  of  oxygen  by  heating  mercuric 
oxide.    How  much  oxide  must  be  heated? 

5.  What  weight  of  hydrogen  will  be  obtained  from  the  action 
of  sulfuric  acid  on  100  g.  of  zinc  ?    What  will  be  its  volume  under 
standard  conditions  ? 

6.  A  given  volume  of  oxygen  standing  over  water  at  20°  and 
745  mm.   measures  10  1.     (a)  What   would   be   its   volume  under 


FORMULAS;  EQUATIONS;  CALCULATIONS    103 

standard  conditions?  (ft)  What  is  its  weight?  (c)  What  weight 
of  potassium  chlorate  would  be  required  to  prepare  this  amount  of 
oxygen  ? 

7.  100  cc.  of  sulfuric  acid  containing  20  g.  of  hydrogen  sulfate 
(H2SO4)  was  added  to  10  g.  of  zinc.    Calculate  the  weight  of  hydro- 
gen evolved. 

8.  WThen  hydrogen  is  liberated  by  the  action  of  zinc  on  sulfuric 
acid,  what  weight  of  zinc  sulfate  is  formed  for  each  gram  of  hydro- 
gen liberated  ? 

9.  If  you  had  10  g.  of  iron  with  which  to  prepare  hydrogen, 
would  you  heat  it  with  steam  or  act  upon  it  with  sulfuric  acid  in 
order  to  obtain  the  maximum  quantity  of  hydrogen  ? 

10.  What  weight  of  potassium  chlorate  is  necessary  to  furnish 
sufficient  oxygen  to  fill  four  200-cubic-centimeter  bottles  in  your 
laboratory  (the  gas  to  be  collected  over  water)  ? 

11.  Calculate  the  weight  of  the  compounds  necessary  for  prepar- 
ing 10  kg.  of  the  common  medicinal  hydrogen  peroxide. 

12.  A  bottle  containing  1  kg.  of  medicinal   hydrogen  peroxide 
was  set  aside  until  the  peroxide  was  completely  decomposed  into 
water  and  oxygen,    (a)  Calculate  the  volume  of  the  oxygen  evolved, 
(ft)  Calculate  the  weight  of  the  water  left  in  the  bottle. 

13.  10  g.  of  zinc  was  used  in  the  preparation  of  hydrogen.    What 
weight  of  iron  will  be  required  to  prepare  an  equal  volume  of  the  gas  ? 

14.  1  kg.  of  potassium  chlorate  was  heated  until  all  the  oxygen 
was  evolved.    Calculate  the  weight  of  the  potassium  chloride  left. 

15.  A  certain  Zeppelin  airship  was  found  to  have  a  capacity  of 
25,000  cu.  yd.    Supposing  that  the  hydrogen  used  in  inflating  this  is 
made  by  the  action  of  sulfuric  acid  on  iron,  calculate  the  weight  of 
hydrogen  sulfate  and  iron  necessary  to  inflate  this  airship  at  20°  and 
740  mm.  (1  yd.  =  0.0144  m.). 


CHAPTER   IX 
THE  THREE  STATES  OF  MATTER 

Gases,  liquids,  and  solids.  We  have  found  that  water 
exists  in  three  very  different  states ;  namely,  as  gas,  as 
liquid,  and  as  solid.  In  a  general  way  these  three  states 
may  be  described  by  saying  that  a  gas  (or  vapor)  is  that 
form  of  matter  that  tends  to  distribute  itself  uniformly 
throughout  the  space  in  which  it  is  placed.  A  liquid 
collects  in  the  bottom  of  the  containing  vessel ;  it  has  no 
characteristic  form  of  its  own,  but  takes  the  shape  of  the 
vessel  in  which  it  is  placed.  A  solid  retains  its  own  form 
irrespective  of  the  shape  of  the  vessel.  Most  substances 
can  be  obtained  in  all  these  states. 

Evaporation.  When  a  liquid  such  as  water  is  placed 
in  an  open  vessel,  it  gradually  passes  into  the  air  in  gase- 
ous form,  or  evaporates.  In  a  confined  space,  as  in  a  partly 
filled  bottle,  evaporation  takes  place  until  the  air  above 
the  liquid  contains  a  definite  percentage  of  vapor,  or  be- 
comes saturated  with  the  vapor.  The  process  of  evapora- 
tion does  not  really  cease  when  saturation  is  reached,  but 
the  rate  at  which  the  vapor  is  formed  from  the  liquid  is 
just  balanced  by  the  rate  at  which  the  vapor  condenses  to 
form  the  liquid.  Saturation  is  therefore  a  balance  between 
these  two  rates. 

If  the  liquid  is  now  warmed,  the  rate  of  evaporation  is 
increased.  A  new  balance  is  reached  at  this  higher  tem- 
perature with  a  higher  percentage  of  vapor  present  in  the 
104 


THE  THREE  STATES  OF  MATTER  105 

air  than  there  was  before,  The  vapor  formed  from  the 
liquid  is  a  gas  and,  like  any  other  gas,  exerts  a  pressure 
upon  the  walls  of  the  containing  vessel  and  upon  the 
surface  of  the  liquid.  This  is  spoken  of  as  the  vapor  pres- 
sure of  the  liquid,  although  it  would  be  more  exact  to  call 
it  the  pressure  of  the  vapor  of  the  liquid.  Solids  as  well 
as  liquids  evaporate,  as  is  shown  by  the  odor  of  such 
substances  as  camphor  and  naphthalene  (moth  balls). 

Relative  humidity.  The  phrase  relative  humidity  is 
familiar  in  the  government  weather  reports.  By  the 
humidity  at  a  given  time  is  meant  the  percentage  of 
water  vapor  in  the  air  as  compared  with  the  percentage 
present  at  saturation  at  the  same  temperature.  To  be 
comfortable,  air  should  be  about  70  per  cent  saturated. 

Boiling  point.  While  a  liquid  is  being  heated,  a  part  of 
the  heat  energy  added  to  it  is  used  in  raising  its  temper- 
ature by  increasing  the  motion  of  the  molecules  of  the 
liquid;  a  part  is  absorbed  in  converting  the  liquid  into 
vapor  against  the  attraction  of  the  molecules  that  tends  to 
hold  it  together  as  a  liquid.  The  formation  of  this  vapor 
is  opposed  by  the  pressure  of  the  atmosphere.  When  the 
pressure  of  the  vapor  just  above  the  liquid  becomes  great 
enough  to  overcome  the  pressure  of  the  atmosphere,  the 
air  is  pushed  back  by  the  vapor.  All  the  heat  energy 
supplied  to  the  liquid  is  now  used  in  changing  the  liquid 
into  vapor  and  in  the  mechanical  work  of  pushing  back 
the  atmosphere,  and  the  temperature  ceases  to  rise.  If 
the  opposing  atmospheric  pressure  is  increased,  the  liquid 
must  be  heated  to  a  higher  temperature  before  its  vapor 
pressure  will  exceed  the  higher  pressure.  The  tempera- 
ture at  which  the  pressure  of  the  vapor  just  exceeds  the 
pressure  of  the  atmosphere  is  called  the  boiling  point  of 
the  liquid.  It  will  be  seen  that  the  point  changes  with 


106    AN  ELEMENTAKY  STUDY  OF  CHEMISTRY 


the  pressure.    Under  a  pressure  of  760  mm.  water  boils  at 
100°;  under  a  pressure  of  525.5  mm.  it  boils  at  90°. 

Heat  of  vaporization  and  of  condensation.  The  quantity 
of  heat  required  to  change  1  g.  of  a  liquid  at  its  boiling 
point  into  1  g.  of  vapor  at  the  same  temperature  is  called 
the  heat  of  vaporization.  For  water  this  is  unusually 
large,  and  amounts  to  539  cal.  If  a  gas  is  maintained 
at  a  pressure  of  760  mm. 
and  is  gradually  cooled,  con- 
densation into  a  liquid  will 
begin  when  the  boiling  point 
is  reached.  During  condensa- 
tion the  temperature  remains 
constant  and  the  quantity  of 
heat  given  out  in  the  process 
(heat  of  condensation)  is  ex- 
actly equal  to  the  heat  of 
vaporization.  Since  the  heat 
given  out  tends  to  oppose  the 
process  taking  place,  conden- 
sation is  not  rapid  unless 
some  method  is  devised  for 


FIG.  40.  The  critical  temperature 
of  a  liquid 


absorbing  the  heat.    This  is  usually  accomplished  by  pass- 
ing the  vapor  through  a  condenser. 

Critical  point.  If  a  liquid  is  sealed  within  a  tube  from 
which  all  air  has  been  withdrawn,  A  (Fig.  40),  the  lower 
end  of  the  tube  will  be  filled  with  liquid  and  the  upper 
end  with  vapor.  If  the  liquid  is  now  heated  by  a 
burner,  it  cannot  boil,  for  the  pressure  of  its  vapor 
cannot  overcome  the  opposing  pressure  and  escape.  As 
the  heating  progresses,  more  and  more  of  the  liquid  is 
vaporized.  The  density  of  the  remaining  liquid  diminishes, 
and  the  density  of  the  vapor  increases.  Evidently  at  some 


THE  THREE  STATES  OF  MATTER  107 

temperature  the  two  will  become  identical,  and  the  bound 
ary  line  between  them,  B  (the  meniscus^),  will  fade  out. 
The  temperature  at  which  this  occurs  is  called  the  critical 
temperature,  and  the  pressure  exerted  by  the  vapor  is 
the  critical  pressure.  Above  this  temperature  no  amount 
of  pressure  will  liquefy  the  gas,  but  will  merely  compress 
it.  Before  any  gas  can  be  liquefied  it  must  first  be  cooled 
below  its  critical  temperature.  The  critical  points  of  a  few 
gases  are  given  in  the  following  table : 

TABLE  OF  CRITICAL  POINTS 

BOILING  .CRITICAL  CRITICAL 

POINT  TEMPERATURE  PRESSURE 


Hydrogen      .     . 

-  252.7° 

-  234.5° 

20.0  atmospheres 

Nitrogen  .     .     . 

-  195.7° 

-  146.0° 

33.0  atmospheres 

Oxygen     .     .     . 

-  182.9° 

-  119.0° 

50.0  atmospheres 

Carbon  dioxide  . 

-  78.2° 

+  31.35° 

72.9  atmospheres 

Water        .     .     . 

+  100.0° 

+  365.0° 

194.6  atmospheres 

Liquefaction  of  gases.  From  what  has  been  said  it  will 
be  clear  that  to  liquefy  a  gas  the  necessary  steps  are 
(1)  to  cool  it  below  its  critical  temperature,  and  (2)  to 
apply  pressure.  At  ordinary  temperatures  many  gases 
are  already  below  their  critical  temperatures  and  so  can 
be  liquefied  by  pressure  alone. 

When  the  critical  temperature  is  very  low,  as  is  true 
with  nitrogen  and  oxygen,  the  temperature  cannot  be 
lowered  sufficiently  by  ordinary  cooling,  but  the  cooling 
is  accomplished  by  mechanical  means.  Machines  con- 
structed for  this  purpose  owe  their  efficiency  to  the  fact 
that  the  heat  given  out  by  a  gas  when  it  is  compressed 
is  not  quite  so  great  as  that  absorbed  when  it  expands. 
Consequently  if  the  gas  is  alternately  compressed  and 
then  allowed  to  expand,  it  will  grow  steadily  colder  and 
may  presently  be  cooled  below  its  critical  temperature. 


108    AN  ELEMENTARY  STUDY  OF  CHEMISTKY 


To  Com 


npressor 


The  Linde  liquid-air  machine.  The  principles  just  explained 
are  well  illustrated  in  the  Linde  machine  for  liquefying  air. 
In  the  Linde  machine  (Fig.  41)  the  compression  is  effected  by 
a  strong  pump.  The  compressed  air  at  200  atmospheres  pres- 
sure is  first  cooled  in  a  freezing  bath  A.  It  then  passes  upward 
as  indicated  by  the  arrow  and  enters  the  inner  tube  of  a  system 
of  three  concentric  spirally  wound  copper  tubes.  At  the  lower 
end  of  this  system  it  expands 
through  a  valve  operated  by  the 
head  screw  B  to  a  pressure  of  from 
20  to  50  atmospheres,  and  in  so 
doing  becomes  much  colder.  It  is 
then  returned  to  the  pump  through 
the  space  between  the  inner  and 
second  tubes  and  the  pipe  at  the 
top,  cooling  the  interior,  com- 
pressed gas.  When  this  process 
no  longer  results  in  a  fall  of  tem- 
perature, the  valve  C  is  opened, 
whereby  some  of  the  cold  air  at 
20  atmospheres  pressure  is  allowed 
to  expand  to  atmospheric  pressure. 

FIG.  41.    The  Linde  machine     In  so  doinS  a  Part  l^efies  and  is 
for  liquefying  air  caught  in  the  vessel  D,  while  the 

very  cold  air  which  escapes  lique- 
faction is  led  back  through  the  outer  tube  of  the  spiral  to  cool 
further  the  air  within  the  two  inner  tubes. 


Dewar  flasks ;  thermos  bottles.  Liquid  air  may  be  kept 
for  some  hours  in  a  special  form  of  flask,  devised  by  the 
Scottish  scientist  Dewar,  known  as  a  Dewar  flask.  This 
consists  of  two  concentric  vessels  (Fig.  42)  of  any  con- 
venient shape.  These  are  joined  together  at  the  upper 
rim  only,  and  the  space  between  them  is  exhausted  by 
an  air  pump.  The  vacuum  serves  as  the  best  possible 
insulator  to  prevent  heat  conduction.  The  surface  of  the 


THE  THREE  STATES  OF  MATTER 


109 


FIG.  42.    A  Dewar 

flask    for    holding 

liquid  air 


outer  flask  is  often  silvered  in  order  to  reflect  the  external 

heat  and  thus  prevent  its  absorption.    The  vessels  known 

as  thermos  bottles  (Fig.  43)  are  constructed  on  the  same  plan 
and  are  very  effective  for  keeping  liquids 
either  hot  or  cold  for  several  hours. 

Solid  bodies.  When  a  liquid  substance 
is  cooled  it  becomes  less  and  less  fluid. 
At  a  sufficiently  low  temperature  all 
liquids  become  rigid  and  we  call  this 
form  a  solid.  As  a  rule  the  change  from 
liquid  to  solid  is  sudden.  At  some  defi- 
nite temperature  crystals  (p.  112)  begin 
to  form  and  the  temperature  of  the  liquid 
comes  to  a  perfectly  definite  value  called 

the  freezing  point,  and  this  remains  unchanged  until  all  of 

the  liquid  has  solidified.    For  example,  water  solidifies,  or 

freezes,  at  0°.    Solids  formed  in  this  way  are 

always  crystalline.     Less  frequently  there 

is   no    definite   point   of  solidification,   the 

changes  from  the  liquid  state  into  that  of 

a  rigid  body  being  very  gradual.    Glasses, 

glazes,  glue,  tar,  and  gums  are  examples 

of  such  materials.     These    are   sometimes 

called  amorphous  solids  to  distinguish  them 

from  crystalline  solids,  but  it  is  better  to 

consider  them  as  still  liquid,  but  so  lacking 

in  fluid  properties  that  they  are  as  rigid  as 

crystalline  solids. 

Freezing   point.     When  a  crystallizable 

liquid,  such  as  water,  is  cooled,  it  does  not 

always  begin  to  crystallize  at  its  freezing  point.    Indeed, 

liquid  water  has  been  cooled  to  —90°  without  freezing.    A 

liquid  below  its  freezing  point  is  said  to  be  undercooled. 


FIG.  43.   A  ther- 
mos bottle 


110    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

If  a  crystal  once  forms  or  one  is  dropped  into  the  under- 
cooled  liquid,  solidification  at  once  begins,  heat  is  given 
out  in  the  process,  and  the  temperature  rises  to  the  true 
freezing  point  and  remains  there  as  long  as  any  liquid  is 
present.  The  freezing  point  is  best  defined  as  the  tem- 
perature at  which  both  solid  and  liquid  will  remain  in 
contact  with  each  other  without  change  of  temperature. 

Heat  of  solidification.  The  heat  given  out  when  1  g.  of  a 
liquid  at  its  freezing  point  solidifies  to  1  g.  of  solid  is  called 
the  heat  of  solidification.  For  water  this  amounts  to  80  cal. 
If  it  were  not  for  this  liberation  of  heat,  which  opposes 
solidification,  ponds  would  freeze  solid  in  the  winter  as 
soon  as  crystallization  had  begun.  Liquids  which  do  not 
have  definite  freezing  points  have  no  heat  of  solidifica- 
tion; this  shows  that  the  rigid  bodies  formed  from  them 
are  fundamentally  different  from  crystalline  solids. 

Melting  point.  If  a  crystalline  solid  is  slowly  heated, 
its  temperature  steadily  rises  to  its  freezing  point  and  the 
change  to  the  liquid  then  begins ;  and  it  has  not  been 
found  possible  to  heat  the  solid  beyond  this  point.  The 
melting  point  and  the  freezing  point  are  therefore  at  the 
same  temperature.  To  convert  1  g.  of  a  solid  at  its  melt- 
ing point  into  a  liquid  at  the  same  temperature  absorbs 
the  same  quantity  of  heat  as  was  liberated  during  solidifi- 
cation. This  is  called  the  heat  effusion.  Amorphous  rigid 
bodies  have  no  definite  melting  point  or  heat  of  fusion. 

The  manufacture  of  ice.  The  manufacture  of  ice  is  based 
on  the  principle  that  in  the  process  of  vaporizing  a  liquid 
a  great  deal  of  heat  is  absorbed  (heat  of  vaporization). 
If  the  process  is  so  conducted  that  the  liquid  in  vaporizing 
absorbs  this  heat  from  water,  the  temperature  of  the  water 
may  be  brought  to  the  freezing  point,  and  by  still  further 
absorption  of  heat  the  water  may  be  frozen.  It  will  be 


THE  THREE  STATES  OF  MATTER 


111 


recalled  that  for  every  gram  of  water  at  0°  frozen  into 
ice  80  cal.  must  be  absorbed.  The  liquid  chosen  to  be 
vaporized  must  readily  pass  into  a  vapor  at  0°  by  lowering 
the  pressure  upon  it,  and  it  should  have  as  great  a  heat 
of  vaporization  as  possible.  Liquid  ammonia  (heat  of 
vaporization,  330  cal.)  is  the  one  most  frequently  used. 

The  general  method  used  in  the  manufacture  of  artificial 
ice  may  be  understood  by  reference  to  Fig.  44.  Ammonia, 
a  gaseous  compound  formed  when  soft  coal  is  heated  in  the 
absence  of  air  (p.  307), 
is  liquefied  by  means 
of  a  compressor  pump 
and  led  into  the  pipes 
A,  B.  The  heat  of 
condensation  is  ab- 
sorbed by  water  flow- 
ing over  the  pipes. 
These  pipes  lead  into 
coils  in  a  large  tank 
nearly  filled  with 
brine,  prepared  by 
dissolving  calcium 
chloride  in  water. 
By  means  of  an  ex- 
pansion valve  C  the 

pressure  upon  the  liquid  ammonia  is  diminished  as  it  enters 
the  coils,  and  the  heat  absorbed  by  the  rapid  evaporation  of 
the  liquid  lowers  the  temperature  of  the  brine  below  0°.  Metal 
vessels  D,  E,  F,  filled  with  pure  water,  are  lowered  into  the 
cold  brine  and  left  until  the  water  in  them  is  frozen  into  cakes 
of  ice.  The  gaseous  ammonia  is  led  through  G  back  to  the 
compressor  pump  and  again  liquefied.  In  a  similar  way  the 
temperature  is  kept  low  in  the .  cold-storage  plants,  now  so 
largely  used  for  preserving  food  products  from  decay.  The 
rooms  of  the  plant  are  supplied  with  pipes,  into  which  liquid 
ammonia  is  forced  and  allowed  to  vaporize. 


FIG.  44.   Diagram  illustrating  the  principle  of 
an  ammonia  ice  machine 


112    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Crystals.  Crystals  may  be  obtained  by  cooling  melted 
solids,  by  the  evaporation  of  solutions,  or  by  cooling  the 
vapors  of  solids.  They  are  solids  bounded  \)j  plane  sur- 
faces, and  these  surfaces  are  arranged  in  orderly  fashion 
Avith  reference  to  coordinate  lines  drawn  through  the  crys- 
tal and  called  the  crystal  axes.  Every  crystal  has  therefore 
a  perfectly  definite  form  (Fig.  45).  Although  there  are 
thousands  of  these  forms,  they  may  all  be  considered  to 


FIG.  45.    Some  typical  examples  of  crystals 

be  modifications  of  six  fundamental  forms,  referred  to  six 
arrangements  of  axes.  A  general  discussion  of  these  forms 
will  be  found  in  the  Appendix. 

Allotropic  forms.  Quite  a  number  of  the  elements  are 
known  to  exist  in  two  or  more  forms  that  are  related  to 
each  other  in  such  a  way  that  one  can  be  converted  into 
the  other  by  the  absorption  or  liberation  of  energy.  •  Such 
modifications  of  an  element  are  called  allotropic  forms. 
Graphite  and  the  diamond  are  two  solid  allotropic  forms 
of  the  element  carbon  (p.  116). 

The  change  of  an  element  into  an  allotropic  modification 
is  very  similar  to  the  change  of  a  solid  into  a  liquid  or 
of  a  liquid  into  a  gas.  The  heat  absorbed  corresponds  to 
the  heat  of  fusion  or  of  vaporization,  and  the  form  that 
has  the  greater  energy  is  the  most  active  chemically,  just 
as  steam  is  more  active  than  water,  and  water  than  ice. 


THE  THREE  STATES  OF  MATTER 


113 


One  of  the  most  interesting  instances  of  an  element 
existing  in  two  allotropic  forms,  and  the  one  earliest 
known,  is  that  of  oxygen  and  ozone. 

Ozone.  As  early  as  1785  the  Dutch  chemist  Van  Marum 
noticed  the  peculiar  odor  that  is  often  observed  near  an 
electrical  machine  when  it  is  discharging  sparks  through 
the  air.  As  the  result  of  a  great 
deal  of  investigation  it  has  been 
shown  that  this  odor  is  due  to  a 
definite  substance  called  ozone. 
Under  ordinary  conditions  ozone 
is  a  pale-blue  gas,  and  when  lique- 
fied it  is  deep-blue  in  color  and 
boils  at  —  119°.  It  can  be  made 
from  pure  oxygen  by  the  action 
of  electrical  discharge,  and  conse- 
quently contains  no  element  other 
than  oxygen. 

Preparation  of-  ozone.  Ozone  is 
most  easily  prepared  by  passing 
a  silent  electric  discharge  (or  elec- 
tric waves)  through  oxygen.  This 
is  done  in  an  apparatus  represented 
in  Fig.  46.  Oxygen  enters  at  A 
and  follows  the  course  indicated  by  the  arrows.  The  metal 
surfaces  B  and  C  are  separated  from  each  other  by  a  space 
through  which  the  oxygen  passes,  and  are  further  insu- 
lated by  the  glass  D.  Wires  from  an  induction  coil  are 
connected  with  B  and  C.  As  the  oxygen  passes  upward 
between  the  metal  plates  it  is  subjected  to  the  electric  dis- 
charge, and  a  portion  of  the  oxygen  is  changed  into  ozone. 

Ozone  is  also  formed  in  many  cases  of  slow  combustion, 
as  when  a  stick  of  phosphorus,  partly  covered  with  water 


FIG.  46.  A  convenient  form 
of  an  apparatus  for  chang- 
ing oxygen  into  ozone 


114    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

to  keep  it  from  taking  fire,  is  exposed  to  air  (Fig.  47). 
It  is  doubtless  formed  during  lightning  discharges,  but  it 
is  doubtful  whether  any  of  it  is  normally  present  in  ah-, 
for  it  is  very  easily  decomposed  into  ordinary  oxygen. 

Conduct  and  uses  of  ozone.    Ozone  resembles  oxygen  in 
many  respects,  but  it  is  very  much  more  active.    Conse- 
quently it  is  a  powerful  oxidizing  agent.     It  acts  upon 
many  colored  substances  (dyes),  con- 
verting them  into  colorless  compounds, 
and  on  this  account  is  used  to  bleach 
many  substances,  such  as  waxes  and 
oils.    It  destroys  many  low  forms  of 
life  and  is  used. as  a  disinfectant  and  for 
purifying  water  for  drinking  supplies. 
The  energy  of  ozone.    Ozone  is  made 
from    oxygen   under   conditions   that 
add  energy  to  the  oxygen.     This  en- 
«W   »  stored   up    in   the   ozone   as 
moist  air  on  phosphorus    chemical  energy  and  is  available  when 
ozone  acts  upon  other  substances.    It 

will  later  be  shown  that  the  molecule  of  ozone  consists  of 
three  oxygen  atoms,  while  that  of  oxygen  consists  of  but 
two.  Ozone  therefore  differs  from  oxygen  (1)  in  chemical 
energy  and  (2)  in  molecular  structure.  The  conversion  of 
oxygen  into  ozone  may  be  represented  in  the  equation 

3  Q2  +  (3  x  21,500  cal.) >•  2  O3 


EXERCISES 

1.  What  are  the  conditions  which  determine  the  physical  state 
of  any  element  ?   Illustrate  in  the  case  of  oxygen. 

2.  Can  all  solid  substances  be  melted  without  decomposition  ? 
Illustrate  by  an  example. 


THE  THREE  STATES  OF  MATTER  115 

3.  Can  all  liquids  be  boiled  without  decomposition?    Illustrate 
by  an  example. 

4.  When  a  pond  begins  to  freeze  in  winter,  why  does  not  all  the 
water  freeze  ? 

5.  Why  does  a  block  of  ice  melt  so  slowly  even  in  warm  air? 

6.  What  becomes  of  the  heat  applied  to  a  boiling  liquid? 

7.  Why  is  it  necessary  to  boil  eggs  longer  on  a  mountain  top 
than  at  the  seashore  in  order  to  cook  them  ? 

8.  Name  three  crystalline  substances. 

9.  How  can  you  crystallize  common  salt? 

10.  Tubs  of  water  are  sometimes  placed  in  cellars,  in  order  to 
prevent  the  free/ing  of  the  stored  fruits  and  vegetables.    Is  this 
practice  based  on  scientific  grounds? 

11.  Suppose  ice  and  water  to  be  mixed  together  at  0°.    Under 
what  conditions  will  more  water  freeze?  more  ice  melt? 

12.  Suggest  a  method  of  raising  the  boiling  point  of  water  above 
100°. 

13.  When  water  freezes  in  a  bottle,  why  is  the  bottle  broken? 
Would  all  other  liquids  act  in  the  same  way  ? 

14.  How  many  calories  of  heat  are  given  off  in  the  freezing  of 
500  g.  of  water  at  0°? 

15.  How  many  calories  are  required  to  change  1  kg.  of  ice  at  0° 
to  water  at  70°? 

16.  How  many  calories  are  required  to  change  1  kg.  of  water  at 
the  temperature  of  your  room  into  steam  at  100°  ? 

17.  What  weight  of  ice  could  be  melted  by  the  heat  evolved  in 
the  condensation  of  50  Ib.  of  steam  at  100°  to  water  at  the  same 
temperature  ?   (1  Ib.  =  453.6  g.) 

18.  100  Ib.  of  ice  at  0°  was  placed  in  a  refrigerator.    The  water 
resulting  from  the  melted  ice  absorbed  sufficient  heat  to  raise  its 
temperature  to  8°  before  it  flowed  from  the  refrigerator.    Calculate 
the  total  number  of  calories  of  heat  absorbed  by  the  ice  and  the 
resulting  water. 


CHAPTER  X 
CARBON  AND  CARBON  DIOXIDE 

Introduction.  Carbon  is  one  of  the  most  familiar  of  the 
elements,  being  present  in  coal  and  charcoal  in  the  free 
state.  Its  most  common  oxide,  called  carbon  dioxide,  is 
formed  in  the  processes  of  respiration  and  combustion, 
and  is  of  great  importance.  It  is  therefore  desirable  that 
we  should  learn  something  of  the  properties  and  chemi- 
cal conduct  of  these  two  substances  at  an  early  stage  in 
our  study. 

Occurrence  of  carbon.  In  the  uncombined  state  carbon 
is  found  in  nature  in  several  forms.  The  diamond  is  virtu- 
ally pure  carbon,  while  graphite  and  the  different  vari- 
eties of  coal  all  contain  more  or  less  free  carbon.  The 
element  also  occurs  abundantly  in  the  form  of  compounds. 
Carbon  dioxide  is  its  most  familiar  gaseous  compound. 
Natural  gas  and  petroleum  are  largely  compounds  of  car- 
bon and  hydrogen.  The  carbonates,  especially  calcium 
carbonate  (limestone),  constitute  great  strata  of  rocks, 
and  are  found  in  almost  every  locality.  All  living  organ- 
isms, both  plant  and  animal,  contain  a  large  percentage  of 
this  element,  and  the  number  of  its  compounds  which  help 
to  make  up  the  great  variety  of  animate  nature  is  almost 
limitless.  In  the  free  state  carbon  occurs  in  both  the 
crystalline  and  the  amorphous  form. 

Crystalline    carbon.    Crystalline   carbon  occurs  in   two 
forms,  —  the  diamond  and  graphite. 
116 


CARBON  AND  CARBON  DIOXIDE 


117 


FIG.  48.   The  Cullinan  diamond  in  its 
original  condition  (one  half  natural  size) 


1.  Diamond.  Diamonds  are  found  in  certain  localities  in 
South  Africa,  the  East  Indies,  and  Brazil.  The  crystals  as 
found  are  usually  covered  with  a  rough  coating.  These  are 
cut  so  as  to  bring  out  the  brilliancy  of  the  gem.  Diamond 

cutting  is  carried  on  most 
extensively  in  Holland. 

The  weight  of  the  dia- 
mond is  usually  expressed 
in  carats,  a  carat  being 
equal  to  about  0.2  g.  The 
word  carat  is  derived  from 
a  Greek  word  meaning  "  the 
seed,  or  bean,  of  the  carob, 
or  locust,  tree."  The  beans 
were  formerly  used  in 
weighing  diamonds. 
The  largest  diamond  known  was  found  in  the  Transvaal 

mines  in  1905,  and  weighed  3025|  carats.    This  was  called  the 

Cullinan  diamond  (Fig.  48)  and  was  presented  to  King  Edward 

VII  by  the  Transvaal  government. 

It  was  subsequently  cut  into  nine 

large   stones    and    a   number    of 

smaller  ones.    The  two  largest  of 

these    weigh    516.5    and    309t\ 

carats  and  are  the  largest  cut  dia- 
monds in  existence.   Other  famous 

diamonds  are  the  Kohinoor  (106J 

carats)  (Fig.  49),  the  Nizam  (277 

carats),  the  Victoria  (180  carats), 

and  the  Jubilee  (239  carats). 

The  density  of  the  diamond 
is  3.5,  and,  though  brittle,  it  is 
one  of  the  hardest  of  substances.  Few  chemical  reagents 
have  any  action  on  it,  but  when  heated  in  oxygen  or  the 
air,  it  blackens  and  burns,  forming  carbon  dioxide. 


FIG.  49.     The     Kohinoor 

diamond   after  being  cut 

(natural  size) 


118    AN  ELEMENTARY  STUDY  OF  CHEMISTKY 

Artificial  production  of  diamonds.  Many  attempts  have  been 
made  to  produce  diamonds  artificially.  For  a  long  time  these 
ended  in  failure,  graphite  and  not  diamonds  being  the  product 
obtained,  but  in  1893  the  French  chemist  Moissan  (Fig.  103), 
in  his  study  of  chemistry  at  high  temperatures,  finally  suc- 
ceeded in  making  some  small  ones.  He  accomplished  this  by 
dissolving  carbon  in  melted  iron  and  plunging  the  crucible 
containing  the  solution  into  water,  as  shown  in  Fig.  50.  Under 
these  conditions  the  carbon  crystallized  in  the  iron  in  the  form 
of  the  diamond.  The  diamonds 
were  then  freed  from  the  metal  by 

(I    V  \/N\rf A,  f**/*'      dissolving  away  the  iron  in  hydro- 
chloric acid. 

2.  Graphite.  This  form  of  car- 
bon is  found  in  large  quantities, 
especially  in  Ceylon,  Siberia,  and 
in  some  parts  of  the  United 
States  and  Canada.  Large  quan- 
—  tities  are  also  made  commercially 

FIG.  50.    Sketch  illustrating    by  heating  hard  coal  to  a  high 
the  method  of  producing  dia-     temperature.     It  is    a  glistening 
monds  in  the  laboratory 

black  substance,  very  soft,  and 

greasy  to  the  touch.  Its  density  is  about  2.15.  It  is  used 
in  the  manufacture  of  lead  pencils  and  crucibles,  as  a 
lubricant,  and,  in  the  form  of  a  polish  or  a  paint,  as 
a  protective  covering  for  iron. 

Commercial  production  of  graphite.  The  process  consists  in  heat- 
ing hard  coal  in  large  electric  furnaces  about  40  ft.  in  length, 
a  longitudinal  section  of  one  of  which  is  shown  in  Fig.  51. 
The  electrodes  A  are  made  of  graphite.  The  furnace  is  nearly 
filled  with  the  coarse  grains  of  coal  (-B).  Since  the  coal  is  a  poor 
conductor,  there  is  placed  in  the  center  of  the  charge  a  core  (C) 
of  carbon,  which  serves  to  conduct  the  current  through  the 
charge.  The  charge  is  covered  with  a  mixture  (D)  of  sand 


CARBON  AND  CARBON  DIOXIDE  119 

and  carbon  (or  similar  materials),  which  excludes  the  air.  An 
alternating  current  is  supplied  by  the  generator  G.  Under  the 
influence  of  the  intense  heat  produced  by  the  current,  the  carbon 
is  changed  into  graphite.  Prepared  in  this  way,  the  product  is 
uniform  in  composition  and  free  from  grit,  and  is  therefore 
superior  to  the  natural  product. 

Amorphous  carbon.  Pure  amorphous  carbon  is  best  pre- 
pared by  heating  sugar  (C12H22O11)  in  the  absence  of  air. 
The  hydrogen  and  oxygen  present  are  expelled,  largely 


FIG.  51.    Electric  furnace  for  the  production  of  graphite 

in  the  form  of  water,  and  pure  carbon  remains.  Among 
the  numerous  substances  that  contain  amorphous  carbon, 
the  following  may  be  mentioned : 

1.  Coal  and  coke.  The  various  forms  of  coal  were  formed 
from  vast  accumulations  of  vegetable  matter.  In  hard,  or 
anthracite,  coal  nearly  all  of  the  carbon  is  in  the  uncom- 
bined  state ;  while  in  soft,  or  bituminous,  coal  a  consid- 
erable portion  of  the  carbon  is  combined  with  hydrogen, 
oxygen,  nitrogen,  and  sulfur.  When  soft  coal  is  heated 
in  the  absence  of  air  (p.  306)  complex  changes  occur, 
resulting  in  the  formation  of  various  useful  compounds 
of  carbon,  which  are  given  off  in  the  form  of  gases  and 
vapors,  while  the  mineral  matter  and  free  carbon  remain 
and  constitute  ordinary  coke.  The  matter  which  escapes 
when  coal  is  heated  in  the  absence  of  air  is  known  as 


120    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


volatile  matter.  In  hard  coal  the  volatile  matter  averages 
from  5  per  cent  to  8  per  cent,  while  in  soft  coal  it  averages 
from  30  per  cent  to  35  per  cent.  When  coal  is  burned,  the 
mineral  matter  present  is  left  in  the  form  of  ash. 

2.  Charcoal.  This  is  prepared  from  wood  just  as  coke 
is  prepared  from  coal.  The  volatile  matter  expelled  con- 
sists of  many  valuable  substances,  such  as  wood  alcohol 


Q 


FIG.  52.   The  modern  method  for  the  production  of  charcoal 

and  acetic  acid,  which  are  obtained  commercially  in  this 
way.  Formerly  much  of  this  volatile  matter  was  allowed 
to  escape,  but  at  present  an  increasing  amount  of  char- 
coal is  prepared  in  such  a  way  that  the  volatile  matter  is 
condensed  and  saved.  Both  charcoal  arid  coke  are  used 
as  fuels,  and  they  are  especially  useful  in  reducing  metals 
from  their  oxides,  as  will  be  described  later. 

Modern  methods  for  the  production  of  charcoal.  Iron  cars  are 
loaded  with  wood,  A,  A  (Fig.  52),  and  run  into  the  retort  B.  The 
retort  is  then  made  air-tight  and  heated  slowly  for  twenty-four 


CARBON  AND  CARBON  DIOXIDE  121 

hours  by  the  fires  F,  F.  The  gaseous  products  escape  through 
the  pipes  C,  C  and  then  pass  into  the  condensers  D,  D.  Here 
those  portions  which  are  liquid  at  ordinary  temperatures,  such 
as  wood  alcohol  and  acetic  acid,  are  condensed,  while  the  gase- 
ous products  are  led  back  into  the  furnace  and  burned.  When 
all  the  volatile  matter  has  been  expelled  in  this  way,  the  cars 
containing  the  charcoal  are  run  into  cooling  chambers,  and  their 
place  in  the  retort  is  taken  by  other  cars  loaded  with  wood. 

3.  Bone  black,  or  animal  charcoal.   This  is  made  by  charring 
bones  and  animal  refuse.    It  consists  of  very  finely  divided 
carbon  and  of  calcium  phosphate,  and  is  especially  useful 
for  removing  coloring  matter  in  the  refining  of  sugar. 

4.  Lampblack.   Lampblack  is  a  product  of  the  imperfect 
combustion  of  carbonaceous  fuels,  such  as  oil  and  gas. 

Destructive  distillation.  The  process  of  decomposing 
such  substances  as  coal,  wood,  and  bones  by  heating  them 
in  the  absence  of  air  is  termed  destructive  distillation. 
Thus,  we  say  that  coke,  charcoal,  and  bone  black  are 
made  by  the  destructive  distillation  of  coal,  wood,  and 
bones,  respectively. 

Properties  of  carbon.  While  the  various  forms  of  carbon 
differ  in  many  properties,  especially  in  color  and  hardness, 
yet  they  are  all  odorless,  tasteless  solids,  insoluble  in  water. 
Only  in  the  intense  heat  of  the  electric  arc  does  carbon 
volatilize,  passing  directly  from  the  solid  state  into  a  vapor 
without  melting.  Owing  to  this  fact,  the  inside  surface  of 
an  incandescent-light  bulb  after  being  used  for  some  time 
becomes  coated  with  a  dark  film  of  carbon.  In  the  form  of 
bone  black,  or  charcoal,  carbon  has  the  property  of  absorbing 
relatively  large  volumes  of  certain  gases,  as  well  as  many 
kinds  of  organic  matter,  from  their  solutions.  As  a  result 
of  this  property,  filtration  through  a  charcoal  filter  will 
often  remove  objectipnable  odors  and  colors  from  solutions. 


122    AN  ELEMENTARY  STUDY  OF  CHEMISTEY 

Chemical  conduct.  At  ordinary  temperatures  carbon  is 
a  very  inactive  substance,  but  at  higher  temperatures  it 
combines  directly  with  a  number  of  elements,  such  as  oxy- 
gen, hydrogen,  and  sulfur.  Because  of  its  strong  affinity 
for  oxygen  at  high  temperatures  it  is  an  excellent  reducing 
agent.  Carbon  also  combines  directly  with  many  of  the 
metals,  forming  compounds  called  carbides.  One  of  the 
most  important  of  these  is  calcium  carbide  (CaC2),  used 
so  largely  in  the  preparation  of  acetylene  and  fertilizers. 
When  heated  in  the  presence  of  sufficient  oxygen,  carbon 
burns,  forming  carbon  dioxide. 

Uses  of  carbon.  The  chief  use  of  amorphous  carbon  is 
for  fuel,  to  furnish  heat  and  power  for  all  the  uses  of 
civilization.  An  enormous  quantity  of  carbon,  in  the  form 
of  coal,  coke,  and  charcoal,  is  used  as  a  reducing  agent 
in  the  separation  of  the  various  metals  from  their  ores. 
Lampblack  is  used  for  making  indelible  ink,  printer's  ink, 
paints,  and  black  varnishes,  while  bone  black  and  charcoal 
are  used  in  water  filters.  In  the  refining  of  sugar  the  dark 
solution  of  the  impure  compound  is  filtered  through  bone 
black,  which  removes  the  coloring  matter.  On  evaporation 
the  resulting  solution  yields  the  colorless  sugar. 

Carbon  dioxide  (C02).  Carbon  dioxide  is  a  colorless  gas 
which  is  formed  whenever  carbon  burns  in  an  atmosphere 
containing  oxygen.  Since  all  of  the  common  fuels,  such 
as  coal,  wood,  oil,  and  gas,  contain  carbon  either  in  a  free 
or  in  a  combined  state,  it  follows  that  carbon  dioxide  is 
formed  whenever  these  fuels  are  burned.  The  gas  is  also 
formed  in  the  process  of  respiration,  from  4  to  5  per  cent 
of  it  being  present  in  exhaled  air.  It  is  likewise  formed 
in  the  processes  of  fermentation  which  take  place  in  the 
manufacture  of  alcohol  and  of  alcoholic  liquors,  as  well 
as  in  making  lime  by  heating  limestone  (p.  425).  Large 


CARBON  AND  CARBON  DIOXIDE 


123 


quantities  of  it  escape  from  volcanoes  and  from  crevices 
in  the  earth.  It  is  present  in  the  air  to  the  extent  of 
about  3  parts  in  10,000,  and  this  apparently  small  quan- 
tity is  of  fundamental  importance  in  nature,  as  will  be 
pointed  out  in  describing  the 
atmosphere. 

Preparation.  In  the  laboratory, 
carbon  dioxide  is  prepared  by 
the  action  of  hydrochloric  acid  on 
the  compound  known  as  calcium 
carbonate  (CaCOg).  The  latter  is 
found  in  nature  in  many  different 
substances,  such  as  shells,  coral, 
and  limestone.  Marble  is  nearly 
pure  calcium  carbonate,  and,  being 
comparatively  inexpensive,  is  the 
material  most  often  used  in  the 
preparation  of  carbon  dioxide.  When  hydrochloric  acid 
and  marble  are  brought  in  contact  with  each  other,  water, 
calcium  chloride  (CaCl2),  and  carbon  dioxide  are  formed 
according  to  the  following  equation: 


FIG.  53.   A  simple  apparatus 
for  preparing  carbon  dioxide 


CaCO3  +  2  HC1 


CaCl2+H20-f-C02 


The  calcium  chloride  is  a  white  solid  which  remains  in 
solution,  while  the  insoluble  carbon  dioxide  escapes  and 
is  collected  by  the  displacement  of  air  as  described  below. 

The  gas  may  be  prepared  in  the  apparatus  shown  in  Fig.  53. 
Pieces  of  marble  are  placed  in  flask  A,  and  hydrochloric  acid 
diluted  with  an  equal  volume  of  water  is  slowly  added  through 
the  funnel  tube  B.  The  tube  C,  through  which  the  carbon 
dioxide  escapes  as  fast  as  formed,  passes  through  a  piece  of 
cardboard  placed  over  the  mouth  of  a  bottle  or  cylinder  as 
shown  in  the  figure.  The  gas  is  heavier  than  air  and  gradually 


124    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


fills  the  cylinder,  pushing  out  the  air.  A  lighted  splint  brought 
into  the  mouth  of  the  cylinder  is  extinguished  when  the 
cylinder  is  filled  with  the  gas. 

The  Kipp  apparatus  (Fig.  17)  is  more  convenient  than  that 
shown  in  Fig.  53.  It  is  used  as  in  the  preparation  of  hydrogen, 
marble  being  substituted  for  the  zinc. 

Properties.  Carbon  dioxide  is  a  colorless,  almost  odor- 
less gas,  one  liter  of  which  weighs  about  1.5  times  as  much 
as  one  liter  of  air.  It  can  be  poured  like  water  from  one 

vessel  downward  into 
another ;  or  as  a  more 
striking  experiment 
the  gas  may  be  poured 
from  a  cylinder  A 
into  a  beaker  J5,  at- 
tached to  a  balance 
and  counterpoised  as 
shown  in  Fig.  54.  At 
15°  and  under  ordi- 
nary pressure  1  vol- 


FIG.  54.    A  method  for  showing  the  weight 
of  gaseous  carbon  dioxide 


ume  of  water  dissolves  1  volume  of  the  gas.  It  is  rather 
easily  condensed  to  a  colorless  liquid,  which  is  slightly 
lighter  than  water  and  boils  at  —  78.2°. 

Liquid  and  solid  carbon  dioxide.  Carbon  dioxide  may  be  pur- 
chased condensed  in  strong  steel  cylinders.  In  these  cylinders 
the  compound  is  under  such  great  pressure  that  it  is  largely 
in  the  liquid  state.  When  the  pressure  is  removed,  the  rapid 
vaporization  of  the  liquid  lowers  the  temperature  sufficiently 
to  freeze  a  portion  of  the  escaping  liquid  to  a  snowlike  solid. 
Cylinders  of  liquid  carbon  dioxide  are  inexpensive  and  should 
be  available  in  every  school.  The  commercial  supply  of  this 
gas  is  obtained  largely  from  fermentation  processes,  especially 
from  breweries. 


CARBON  AND  CARBON  DIOXIDE  125 

To  prepare  the  solid  carbon  dioxide,  the  cylinder  should  be 
placed  across  the  table  and  supported  in  such  a  way  that  the 
stopcock  end  is  several  inches  lower  than  the  other  end.  A 
loose  bag  is  made  by  holding  the  corners  of  a  piece  of  cloth 
around  the  neck  of  the  stopcock.  The  stopcock  is  then  turned 
on  so  that  the  liquid  rushes  out  in  large  quantities.  Very 
quickly  a  considerable  quantity  of  the  snow  collects  in  the  cloth. 
Mercury  may  be  frozen  by  this  snow  in  the  following  way :  A 
filter  paper  is  placed  in  the  bottom  of  a  small  evaporating  dish 
and  some  mercury  poured  upon  it.  One  end  of  a  piece  of  wire 
is  wound  into  a  flat  coil  and  dipped  into  the  mercury.  A  quan- 
tity of  the  solid  carbon  dioxide  is  placed  upon  the  mercury 
and  from  10  to  15  cc.  of  ether  poured  over  it.  The  tempera- 
ture is  reduced  to  —  50°,  so  that  the  mercury  solidifies  in  a 
minute  or  two  and  may  be  removed  from  the  dish  by  the  wire 
which  serves  as  a  handle. 

Chemical  conduct.  Carbon  dioxide  is  a  very  stable  sub- 
stance. It  is  neither  combustible  nor  a  supporter  of  com- 
bustion. When  it  is  passed  into  a  clear  solution  of  calcium 
hydroxide,  Ca(OH)2  (ordinary  limewater),  the  solution 
soon  becomes  cloudy,  owing  to  the  formation  of  calcium 
carbonate.  This  substance  is  insoluble  and  separates  out 
as  fast  as  it  is  formed,  producing  a  cloudy  or  milky 
appearance  in  the  solution: 

Ca(OH)2  +  C02 *  CaC03  +  H2O 

These  properties  constitute  a  simple  test  for  carbon  dioxide. 
Uses.  The  carbon  dioxide  in  the  air  is  a  food  for  plants, 
as  will  be  shown  in  the  chapter  on  the  atmosphere.  Com- 
mercially it  is  used  chiefly  in  the  manufacture  of  soda 
water  and  similar  beverages  and  as  a  fire  extinguisher. 
Ordinary  soda  water  consists  of  different  flavoring  extracts 
to  which  is  added  water  charged  with  carbon  dioxide  under 
pressure.  When  the  pressure  is  removed,  the  excess  of 


126    AN  ELEMENTARY   STUDY  OF  CHEMISTRY 


gas  escapes,  producing  effervescence.  Most  of  the  portable 
fire  extinguishers  are  simply  devices  for  generating  carbon 
dioxide.  It  is  not  necessary  that  all  the  oxygen  should 
be  kept  away  from  a  fire  in  order  to  smother  it.  A  burn- 
ing candle,  for  example,  is  extin- 
guished in  air  which  contains  only 
2.5  per  cent  of  carbon  dioxide. 

Fire  extinguishers.  The  general 
type  of  the  portable  fire  extin- 
guisher is  shown  in  Fig.  55.  The 
liquid  is  a  solution  of  sodium  hy- 
drogen carbonate  in  water.  The 
bottle  A  contains  sulfuric  acid  in 
sufficient  amount  to  react  with  the 
sodium  carbonate  in  solution.  In 
case  of  fire,  the  extinguisher  is 
caught  by  the  handle  D  and  inverted, 
and  the  bottle  containing  the  sul- 
furic acid  is  broken'  by  striking  the 
rod  B  against  the  floor.  The  sul- 
furic acid  immediately  reacts  with 
the  carbonate,  generating  carbon 
dioxide,  some  of  which  dissolves  in  the  water,  while  the 
remainder  forces  the  solution  out  through  the  nozzle  C.  While 
the  total  quantity  of  water  furnished  by  such  an  extinguisher 
is  comparatively  small,  it  is  very  effective  as  a  fire  extinguisher, 
because  of  the  large  percentage  of  carbon  dioxide  which  it 
contains  in  solution. 

EXERCISES 

1.  Suggest  a  method  for  proving  that  all  the  various  forms  of 
carbon  described  are  really  carbon. 

2.  How  could  you  judge  of  the  relative  purity  of  different  forms 
of  carbon? 

3.  How  could  one  distinguish  between  oxygen,  hydrogen,  and 
carbon  dioxide? 


FIG.  65.    An  example  of  a 
modern  fire  extinguisher 


CARBON  AND  CARBON  DIOXIDE  127 

4.  Suggest  a  method  for  determining  the  percentage  of  carbon 
in  a  sample  of  coal. 

5.  Apart  from  its  color,  why  should  carbon  be  useful  in  the 
preparation  of  inks  and  paints  ? 

6.  Report   important   events    in  the    life   of   Moissau  (consult 
encyclopedia). 

7.  («)  Calculate  the  weight  of  100  1.  of  carbon  dioxide,  (b)  What 
weight  of  marble  is  necessary  for  the  preparation  of  this  volume  of 
the  gas  ?    (c)  What  weight  of  calcium  chloride  would  be  formed  in 
this  process  ? 

8.  Contrast  the  boiling  points  of  carbon  dioxide,  oxygen,  and 
hydrogen. 

9.  What  effect  would  doubling  the  pressure  have  upon  the  solu- 
bility of  carbon  dioxide  in  water  ? 

10.  Enumerate  the  important  products  resulting  from  the  destruc- 
tive distillation  of  wood. 

11.  Why  does  soda  water  effervesce? 

12.  The  reaction  which  takes  place  when  the  sulfuric  acid  and 
sodium    hydrogen  carbonate   in    a   fire    extinguisher   are    brought 
together  is  represented  by  the  following  equation : 

2  NaHCO3  +  H2SO4 >-Na2SO4  +  2  H2O  +  2  CO2 

(«)  What  weight  of  hydrogen  sulfate  would  have  to  be  used  for 
each  kilogram  of  the  carbonate  dissolved  in  the  water?  (6)  What 
volume  of  carbon  dioxide  would  be  evolved  ? 

13.  How  could  you  prove  that  carbon  dioxide  is  a  product  of 
combustion  (p.  3)  ? 

14.  How  could  you  convert  the  carbon  dioxide  in  the  air  which 
you  exhale,  into  calcium  carbonate  ? 

15.  What  weight  of  calcium  carbonate  would  be  necessary  to 
prepare  sufficient  carbon  dioxide  to  saturate  10  1.  of  water  at  15°  and 
under  ordinary  pressure  ? 

16.  Suppose  the  Kohinoor  diamond  were  to  be  burned  in  oxygen, 
calculate  the  volume  of  the  product  of  combustion. 


CHAPTER  XI 


NITROGEN  AND  THE  RARE  ELEMENTS  IN  THE 
ATMOSPHERE 

Historical.  Nitrogen  was  discovered  by  the  Scottish 
chemist  Rutherford  in  1772.  A  little  later  Scheele  (Fig.  56) 
showed  it  to  be  a  constituent  of  air,  and  Lavoisier  gave  it 

the  name  azote,  signifying 
that  it  would  not  support 
life.  The  name  nitrogen  was 
afterwards  suggested  be- 
cause of  its  presence  in  salt- 
peter, or  niter.  The  term 
azote  and  the  symbol  Az  are 
still  used  by  the  French 
chemists. 

Occurrence.  Air  is  com- 
posed principally  of  oxygen 
and  nitrogen,  each  in  the 
free  state  —  about  78  parts 
by  volume  out  of  every  100 
parts  being  nitrogen.  Nitro- 
gen also  occurs  in  nature  in 
potassium  nitrate  (KNO3) 
— commonly  called  saltpeter 
or  niter  —  as  well  as  in  sodium  nitrate  (NaNOg)  (Fig.  163). 
Nitrogen  is  also  an  essential  constituent  of  all  living  organ- 
isms ;  for  example,  the  human  body  contains  about  3  per 
cent  of  nitrogen. 

128 


FIG.  56.    Karl  Wilhelm  Scheele 
(1735-1784) 

A  Swedish  chemist  who  made  many 
discoveries  in  chemistry 


NITKOGEN  AND  THE  RAKE  ELEMENTS      129 


Preparation  from  air.  Nitrogen  can  be  prepared  from 
air  by  the  action  of  some  substance  which  will  combine  with 
the  oxygen,  leaving  the  nitrogen  free.  Such  a  substance 
must  be  chosen,  however,  as  will  combine  with  the  oxygen 
to  form  a  product  which  is  not  a  gas  and  which  can  be 
readily  separated  from  the  nitrogen.  The  substances  most 
often  used  for  this  purpose  are  phosphorus  and  copper. 

1.  By  the  action   of  phosphorus.    The  method  used  for 
the  preparation  of  nitrogen  by  the  use  of  phosphorus  is 
as  follows: 

The  phosphorus  is  placed  in  a  little 
porcelain  dish  supported  on  a  cork  and 
floated  on  water  (Fig.  57).  It  is  then 
ignited  by  contact  with  a  hot  wire, 
and  immediately  a  bell  jar  or  bottle 
is  brought  over  it  so  as  to  confine  a 
portion  of  the  air.  The  phosphorus 
combines  with  the  oxygen  to  form  an 
oxide  of  phosphorus  known  as  phos- 
phorus pentoxide.  This  is  a  white 
solid  which  floats  about  in  the  bell 

jar,  but  which  in  a  short  time  is  all  absorbed  by  the  water, 
leaving  the  nitrogen.  The  withdrawal  of  the  oxygen  is  indi- 
cated by  the  rising  of  the  water  in  the  bell  jar. 

2.  By  the  action  of  copper.    The  oxygen  in  the  air  may 
also  be  removed  by  passing  air  slowly  through  a  heated 
tube  containing  copper.    The   copper  combines  with  the 
oxygen    to   form   copper   oxide,    which   is   a   solid.     The 
nitrogen    passes    on    and    may    be    collected    over    water. 
The  details  of  the  process  are  as  follows: 

The  copper  is  placed  in  a  tube  A  (Fig.  58)  and  heated.  Air 
is  then  forced  slowly  through  the  tube  by  pouring  water  into 
the  bottle  B.  The  oxygen  of  the  air  combines  with  the  hot 


FIG.  57.    Preparing  ni- 
trogen by  burning  out 
the  oxygen  of  air  with 
phosphorus 


130    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

copper,  forming  the  black  solid,  copper  oxide  (CuO),  which 
remains  in  the  tube,  while  the  nitrogen  passes  on  and  is 
collected  over  water  in  the  cylinder  C. 

The  nitrogen  obtained  from  air  by  the  above  methods 
is  never  quite  pure,  but  contains  about  1  per  cent  of  a 
mixture  of  other  gases  (p.  133).  The  properties  of  the 
nitrogen,  however,  are  not  materially  affected  by  the  pres- 
ence of  these  gases.  If  we  wish  to  obtain  pure  nitrogen, 
we  may  do  so  by  heating  certain  compounds  of  nitrogen. 


FIG.  58.   Preparing  nitrogen  by  removing  the  oxygen  from  ah- 
with  hot  copper 

Preparation  from  compounds  of  nitrogen.  The  compound 
most  often  used  for  the  preparation  of  nitrogen  is  ammonium 
nitrite  (NH4NO2).  When  heated,  this  compound  decom- 
poses into  nitrogen  and  water,  as  represented  in  the  follow- 
ing equation : 

NH4N02— ^N2  +  2H20 

Since  ammonium  nitrite  is  not  readily  kept  in  the  pure  state, 
it  is  convenient  to  substitute  for  it  a  mixture  of  sodium  nitrite 
(NaNO2)  and  ammonium  chloride  (NH4C1).  These  two  com- 
pounds react  to  form  sodium  chloride  and  ammonium  nitrite : 

NaNO,  -(-  NH4C1 >•  NH4N02  +  NaCl 


NITROGEN  AND  THE  RARE  ELEMENTS      131 

Commercial  methods  of  preparation.  It  is  evident  that 
the  method  used  for  the  preparation  of  oxygen  from  liquid 
air  (p.  28)  would  serve  equally  well  for  the  preparation 
of  nitrogen.  This  method  has  come  to  the  front  during  the 
World  War  through  the  development  of  more  efficient 
machines  for  liquefying  air  (p.  108).  By  this  method  is 
prepared  the  large  amount  of  nitrogen  used  in  the  manu- 
facture of  ammonia  by  the  Haber  Process  (p.  203)  as  well 
as  of  certain  fertilizers. 

Properties.  Nitrogen,  like  hydrogen  and  oxygen,  is  a 
colorless,  odorless,  tasteless  gas.  It  is  slightly  lighter  than 
oxygen,  1  1.  of  it  weighing  1.2507  g.  Its  solubility  in 
water  is  about  the  same  as  that  of  hydrogen,  1  1.  of 
water  dissolving  about  20  cc.  of  the  gas  under  standard 
conditions.  Liquid  nitrogen  is  colorless,  boils  at  —195.7°, 
and  has  a  density  of  0.8  at  its  boiling  point.  At  a  still 
lower  temperature  it  may  be  obtained  in  the  form  of  an 
icelike  solid  melting  at  —  210.5°. 

Chemical  conduct.  Nitrogen  is  much  less  active  than 
oxygen,  showing  little  or  no  tendency  to  combine  with 
any  other  elements  at  ordinary  temperatures.  At  higher 
temperatures  it  combines  with  magnesium,  lithium,  titanium, 
and  a  number  of  other  elements.  The  compounds  formed 
are  in  general  called  nitrides,  just  as  compounds  of  an 
element  with  oxygen  are  called  oxides.  When  nitrogen  is 
mixed  with  oxygen  and  subjected  to  the  action  of  electric 
sparks,  the  two  gases  slowly  combine,  and  form  oxides  of 
nitrogen.  A  mixture  of  nitrogen  and  hydrogen  when  treated 
similarly  forms  ammonia  (NH8),  a  gaseous  compound  of 
nitrogen  and  hydrogen.  Since  we  are  constantly  inhaling 
nitrogen,  it  is  evident  that  it  is  not  poisonous.  Nevertheless 
life  would  be  impossible  in  an  atmosphere  of  pure  nitrogen 
on  account  of  the  exclusion  of  the  necessary  oxygen. 


132    AN  ELEMENTAKY  STUDY  OF  CHEMISTRY 


The  assimilation  of  nitrogen  by  plants.  While  nitrogen 
is  an  essential  constituent  of  both  plants  and  animals, 
yet,  with  the  exception  of  a  few  plants,  especially  those 
belonging  to  the  natural  order  Leguminosae,  these  organ- 
isms have  not  the  power  of  directly  assimilating  free 
nitrogen  from  the  atmosphere,  but  obtain  their  supply 
from  certain  compounds  of  nitrogen.  It  has  long  been 
known  that  some  of  the  leguminous  plants,  such  as  the 
beans,  peas,  and  clover,  not  only  thrive  in  poor  soil  but 
at  the  same  time  enrich  it.  Investigation  has  shown  that 
these  plants  obtain  at  least  a  por- 
tion of  their  supply  of  nitrogen 
from  the  atmosphere.  The  as- 
similation of  nitrogen  is  accom- 
plished through  the  agency  of 
groups  of  microorganisms  which 
produce  little  tubercles  on  the 
roots  of  the  plants,  as  repre 
sented  hi  Fig.  59.  The  figure 
shows  the  tubercles  on  the  roots 
of  a  variety  of  bean.  These  microorganisms  have  the 
power  of  converting  free  nitrogen  taken  from  the  air  into 
compounds  of  nitrogen,  some  of  which  are  assimilated 
by  the  plant,  while  others  are  left  in  the  soil  and  thus 
enrich  it. 

Uses  of  nitrogen.  Free  nitrogen  is  used  to  a  limited 
extent  in  the  preparation  of  certain  nitrogenous  compounds 
(p.  462)  employed  as  fertilizers.  Mercurial  thermometers 
designed  for  use  at  temperatures  of  from  300°  to  500°  are 
filled  (over  the  mercury)  with  nitrogen  under  pressure. 
In  this  way  the  mercury  is  prevented  from  boiling,  even 
at  temperatures  considerably  above  its  ordinary  boiling 
point  (357°). 


FIG.  69.  Tubercles  on  the  roots 
of  bean  plants 


NITROGEN  AND  THE  RARE  ELEMENTS     133 


Argon,  helium,  neon,  krypton,  xenon.  These  are  rare  elements 
and  occur  in  the  air  in  very  small  quantities.  They  are  similar 
in  that  they  are  all  colorless,  odorless  gases.  They  differ  from  all 
other  known  elements  in 
that  they  are  entirely  inert, 
forming  no  compounds 
whatever.  Argon,  the  most 
abundant  of  the  group, 
was  discovered  in  1894 
by  two  British  scientists, 
Lord  Rayleigh  and  Sir 
William  Ramsay  (Fig.  60). 
In  1868  Lockyer  showed 
that  a  gaseous  element,  to 
which  he  gave  the  name 
helium,  was  present  in  the 
gases  surrounding  the  sun. 
In  1895  Ramsay  showed 
that  this  same  element 
was  present  in  the  gases 
evolved  in  heating  certain 
minerals,  and  later,  that 
traces  of  it  were  present 
in  the  atmosphere.  The 
three  remaining  members 
of  the  group  were  discov- 
ered in  liquid  air  by  Ramsay  and  Travers  in  1898. 

The  commercial  preparation  oif  helium.  Helium  is  also  present 
to  the  extent  of  nearly  1  per  cent  in  the  natural  gas  found  in 
certain  localities,  especially  in  Kansas  and  Texas.  When  the 
United  States  entered  the  war,  experimental  plants  were  built  in 
Texas  for  the  purpose  of  developing,  if  possible,  a  commercial 
method  for  the  separation  of  the  helium  present  in  natural  gas, 
in  the  hope  that  the  element  could  be  obtained  in  sufficient 
quantities  to  allow  its  use  as  a  material  for  filling  balloons  and 
dirigibles ;  for,  while  helium  is  not  so  light  as  hydrogen,  it 
possesses  a  very  great  advantage  over  hydrogen,  as  a  material 


FIG.  60.    Sir  William  Kamsay 
(1852-1916) 

Famous  for  his  discovery  *.f  the  rare  gases 

in  the  atmosphere  and  far  his  researches 

in  radioactivity 


134     AN  ELEMENTAKY  STUDY  OF  CHEMISTEY 


for  filling  balloons,  in  being  non inflammable.  These  experi- 
ments were  successful,  and,  had  the  war  continued  a  few 
months  longer,  an  adequate  supply  would  have  been  at  hand. 
The  method  used  for  the  separation  of  helium  was  similar 
to  that  used  in  the  preparation  of  oxygen  and  nitrogen  from 
liquid  air,  advantage  being  taken  of  the  fact  that  helium  has  a 
very  low  boiling  point,  namely,  —  268.7°. 
I  Some  facts  pertaining  to  the  rare  gases  in  the  atmosphere 
are  given  in  the  following  table : 


HELIUM 

NEON 

ARGON 

KRYPTON 

XENON 

Weight  of  1  1.  of  gas  . 

0.1  782  g. 

0.9002g. 

1.7809g. 

3.708g. 

5.851  g. 

Boiling  point  of  liquid 

-268.7° 

-239° 

-  186° 

-  151.7° 

-  109° 

Number  of  volumes  in 

1,000,000  volumes  of 

air  (approximate)    . 

4.00 

12.3 

9400 

0.05 

0.006 

1.  How  eould  you  distinguish  between  oxygen,  hydrogen,  and 
nitrogen  ? 

2.  Calculate  the  relative  weights  of  nitrogen  and  oxygen;  of 
nitrogen  and  hydrogen. 

3.  In  the  preparation  of  nitrogen  from  the  air,  how  would  hydro- 
gen do  as  a  substance  for  the  removal  of  the  oxygen  ? 

4.  Why  not  prepare  nitrogen  by  burning  a  candle  in  confined 
air? 

5.  Which  contains  the  greater  percentage  of  nitrogen,  sodium 
nitrate  or  potassium  nitrate  ? 

6.  What  is  the  significance  of  each  of  the  following  names: 
argon,  helium,  neon,  krypton,  xenon  ?    (Consult  dictionary.) 

7.  Note  some  of  the  important  discoveries  made  by  Scheele 
(consult  encyclopedia). 

8.  What  weight  of  nitrogen  can  be  obtained  from  10 1.  of  air 
measured  under  the  conditions  of  temperature  and  pressure  which 
prevail  in  your  laboratory  ? 


CHAPTER  XII 
THE  ATMOSPHERE 

Historical.  The  terms  atmosphere  and  air  are  often  used 
interchangeably,  although  strictly  speaking  the  former  term 
is  applied  to  the  entire  gaseous  envelope  surrounding  the 
earth,  while  the  latter  is  applied  to  a  limited  portion  of 
this  envelope.  Like  water,  air  was  formerly  regarded  as 
an  element.  Near  the  close  of  the  eighteenth  century, 
however,  through  the  experiments  of  Scheele,  Priestley, 
Cavendish,  and  Lavoisier,  it  was  shown  to  be  a  mixture  of 
at  least  two  gases,  —  those  which  we  now  call  oxygen  and 
nitrogen.  By  absorbing  the  oxygen  from  an  inclosed  vol- 
ume of  air  and  measuring  the  contraction  in  volume  caused 
by  the  removal  of  oxygen,  Cavendish  was  able  to  determine 
with  considerable  accuracy  the  relative  volumes  of  oxygen 
and  nitrogen. 

Composition  of  the  air.  The  normal  constituents  of  air, 
together  with  the  approximate  volumes  of  each  in  samples 
collected  in  the  open  fields,  are  as  follows: 

Oxygen 21  volumes  in  100  volumes  of  dry  air 

Nitrogen 78  volumes  in  100  volumes  of  dry  air 

Water  vapor   ....  variable  within  wide  limits 

Carbon  dioxide    .     .     .  3  to  4  volumes  in  10,000  volumes  of  dry  air 

Argon 0.940  volumes  in  100  volumes  of  dry  air 

Helium,  neon,  krypton, 

xenon traces 

In  addition,  there  are  usually  present  small  quantities  of 

hydrogen  peroxide,  oxides  of  nitrogen,  ammonium  nitrate, 

135 


136    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

microorganisms,  dust  particles,  and  traces  of  hydrogen. 
The  air  in  large  cities  and  manufacturing  districts  is  also 
likely  to  contain  certain  gases  evolved  in  manufacturing 
processes.  Among  these  are  hydrogen  sulfide  (H2S)  and 
sulfur  dioxide  (SO2). 

Water  vapor  in  the  air.  The  quantity  of  water  vapor 
which  may  be  present  in  the  air  varies  with  the  tempera- 
ture. This  is  shown  in  the  following  table,  which  gives 
the  weight  in  grams  of  the  water  vapor  that  1  cu.  m.  of 
air  can  absorb  at  the  temperature  indicated : 

Temperature,  0°  10°  20°  30° 

Weight  of  water,         4.8  g.         9.9  g.         17.1  g.         30  g. 

The  constituents  of  the  air  that  are  essential  to  life.  The 
constituents  that  are  known  to  be  essential  to  lif e  are 
oxygen,  nitrogen,  water  vapor,  and  carbon  dioxide. 

The  oxygen  in  the  atmosphere  directly  supports  life 
through  the  process  of  respiration.  The  nitrogen  serves 
to  dilute  the  oxygen  and  thus  to  diminish  the  intensity 
of  its  action.  It  is  likewise  assimilated  by  certain  plants 
(p.  132).  The  water  vapor  prevents  excessive  evaporation 
of  the  water  present  in  organisms,  while  the  carbon  dioxide 
is  an  essential  plant  food. 

The  quantitative  analysis  of  air.  A  number  of  different 
methods  have  been  devised  for  the  determination  of  the 
percentages  of  the  constituents  of  the  atmosphere.  Among 
these  are  the  following: 

1.  Determination  of  oxygen.  The  oxygen  is  withdrawn, 
by  means  of  phosphorus,  from  a  measured  volume  of  air 
inclosed  in  a  tube. 

To  make  the  determination,  a  graduated  tube  is  filled  with 
water  and  inverted  in  a  vessel  of  water.  A  sample  of  the  air 
to  be  analyzed  is  then  introduced  into  the  tube  until  the  tube  is 


THE  ATMOSPHERE 


137 


nearly  filled  with  the  gas,  and  the  volume  is  carefully  noted.  A 
small  piece  of  phosphorus  is  attached  to  a  wire  and  brought 
within  the  tube  as  shown  in  Fig.  61.  After  a  few  hours  the 
oxygen  in  the  inclosed  air  will  have  combined  with  the  phos- 
phorus, the  water  rising  to  take  its  place.  The  phosphorus  is 
removed,  and  the  volume  is  again  noted.  The  contraction  in 
the  volume  of  the  air  is  equal  to  the 
volume  of  oxygen  absorbed  from  the  air. 

2.  Determination  of  nitrogen.   If  the  gas 
left  after  the  removal  of  oxygen  from  a 
portion    of    air   is   passed    over   heated 
magnesium    or    lithium,    the     nitrogen 
present  in  the  gas  is  withdrawn,  leav- 
ing argon  and  the  other  rare  elements. 
It  may  thus  be  shown  that,  of  the  79 
volumes  of  gas  left  after  the  removal 
of  the  oxygen  from  100  volumes  of  dry 
air,  approximately  78  volumes  are  nitro- 
gen and  0.94  argon.   The  other  elements 
are  present  in  such  small  quantities  that 
they  may  be  neglected. 

3.  Determination  of  water  vapor  and  carbon  dioxide.    These 
constituents  are  determined  by  passing  a. known  volume 
of  air  through  two  tubes,  the  first  containing  calcium  chlo- 
ride, and  the  second  calcium  hydroxide,  or,  better,  sodium 
hydroxide.     The  calcium  chloride  removes  the  moisture, 
while  fhe  sodium  hydroxide  removes  the  carbon  dioxide. 
The  increase  in  the  weights  of  these  two  substances  will 
give  the  weights  of  moisture  and  carbon  dioxide  respectively 
in  the  original  volume  of  air. 

Processes  tending  to  change  the  composition  of  the  air. 
These  processes  fall  into  two  classes :  those  which  increase 
the  carbon  dioxide  and  those  which  diminish  it. 


FIG.  61.    The  deter- 
mination of  the  oxy- 
gen in  air  by  means 
of  phosphorus 


138    AX  ELEMENTARY  STUDY  OF  CHEMISTRY 


1.  Processes  tending  to  increase  the  quantity  of  carbon  dioxide. 

Not  only  do  large  quantities  of  carbon  dioxide  escape  into 
the  atmosphere  from  volcanoes  and  crevices  in  the  earth's 
crust,  but  certain  processes   are   constantly  taking  place 
which  are  attended  by  evolution  of  this  gas.    Chief  among 
these  are  the  following:   (a)  Respiration.    In  this  process 
some  of  the  oxygen  in  the  inhaled  air  is  absorbed  by  the 
blood  and  carried  to  all  parts  of 
the  body,  where  it  combines  with 
the  carbon  present  in  the  tissues 
of  the  body,  the  reaction  being 
attended  by   a  transformation 
of  chemical   energy   into   heat 
and  muscular  energy.  The  prod- 
ucts   of    oxidation    are    carried 
back  to  the  lungs  and  exhaled 
largely  in  the  form  of  carbon 
dioxide,    (b)    Combustion.    A-ll 
the  ordinary  fuels  contain  large 
percentages  of  carbon.  On  burn- 
ing, this  is  oxidized  to  carbon 
dioxide,     (c)  Decay  of  organic 
matter.     When  organic  matter 
decays  in  the  air  the  carbon  pres- 
ent is  oxidized  to  carbon  dioxide. 

2.  Processes  tending  to  decrease  the  quantity  of  carbon  dioxide. 
There  are  two  general  processes  which  tend  to  diminish  the 
quantity  of  carbon  dioxide  in  the  atmosphere. 

(a)  The  action  of  plants.  Plants  have  the  power,  when 
growing  in  sunlight,  of  absorbing  carbon  dioxide  from  the 
air,  retaining  the  carbon  and  returning  a  portion  of  the 
oxygen  to  the  air.  It  is  from  this  source  that  plants  obtain 
their  entire  supply  of  carbon. 


FIG.  62.  The  liberation  of  oxy- 
gen   from    plants*  exposed    to 
sunlight 


THE  ATMOSPHERE  139 

That  plants  evolve  oxygen  in  the  sunlight  may  be  shown  as 
follows :  Some  freshly  gathered  leaves  are  placed  under  water 
in  the  jar  A  (Fig.  62)  and  covered  with  the  funnel  B,  the  stem 
of  which  extends  into  the  graduated  tube  C.  Bubbles  of  oxygen 
make  their  escape  from  the  surface  of  the  leaves  and  may  be 
collected  in  the  measuring  tube  C. 

(b)  The  weathering  of  rocks.  Large  quantities  of  carbon 
dioxide  are  being  constantly  withdrawn  from  the  atmos- 
phere through  its  combination  with  various  rock  materials. 

The  composition  of  the  air  constant.  Notwithstanding  the 
changes  taking  place  which  tend  to  alter  the  composition 
of  the  air,  the  results  of  the  analyses  of  air  show  that  the 
percentages  of  oxygen  and  nitrogen,  as  well  as  of  carbon 
dioxide,  are  very  nearly  constant.  Indeed,  so  constant  are 
the  percentages  of  oxygen  and  nitrogen  that  the  question 
has  arisen,  whether  air  is  not  a  definite  chemical  compound. 

Air  a  mixture.  That  the  oxygen  and  nitrogen  in  the 
air  are  not  combined  may  be  shown  in  a  number  of  ways, 
among  which  are  the  following : 

1.  When  air  dissolves  in  water  it  has  been  found  that 
the  ratio  of   oxygen  to  nitrogen   in  the   dissolved  air  is 
no  longer  21 :  78,  but  more  nearly  35  :  65.    If  air  were  a 
chemical  compound,  the  ratio  of  oxygen  to  nitrogen  would 
not  be  changed  by  solution  in  water. 

2.  A  chemical  compound  in  the  form  of  a  liquid  has  a 
definite  boiling  point  at  a  given  pressure  (p.  105).    Water, 
for  example,  boils  at  100°  under  standard  pressure.    More- 
over, the  steam  which  is  formed  has  the  same  composition 
as  the  water.    The  boiling  point  of  liquid  air,  on  the  other 
hand,  gradually  rises  as  the  liquid  boils,  the  nitrogen  es- 
caping first,  then  the  oxygen.     If  the  two  were  combined, 
they  would  pass  off  together  in  the  ratio  in  which  they 
are  found  in  the  air. 


140    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Why  the  air  has  a  constant  composition.  If  air  is  a  mixture 
and  changes  are  constantly  taking  place  which  tend  to  modify 
its  composition,  how,  then,  do  we  account  for  the  constancy  of 
composition  which  the  analyses  reveal  ?  This  is  explained  by 
several  facts :  (1)  The  changes  which  are  caused  by  the  proc- 
esses of  combustion,  respiration,  and  decay,  on  the  one  hand, 
and  the  action  of  plants,  on  the  other,  tend  to  equalize  each 
other.  (2)  The  winds  keep  the  air  in  constant  motion  and  so 
prevent  local  changes.  (3)  The  volume  of  air  is  so  vast  and  the 
changes  which  occur  are  so  small,  compared  with  the  total  vol- 
ume, that  they  cannot  be  readily  detected.  (4)  Finally,  it  must 
be  noted  that  only  air  collected  in  the  open  fields  shows  this 
constancy  in  composition.  The  air  in  a  poorly  ventilated  room 
occupied  by  a  number  of  people  rapidly  changes  in  composition. 

Impure  air  and  ventilation.  The  difference  in  the  per- 
centages of  oxygen,  carbon  dioxide,  and  moisture  present 
in  inhaled  and  exhaled  air  are  shown  in  the  following  table : 


CONSTITUENT 

INHALED  AIR 

EXHALED  AIR 

Oxygen     
Carbon  dioxide  

21.00% 
0.04% 
variable 

16.00% 
4.38% 
saturated 

The  injurious  effects  resulting  from  inadequate  ventilation 
seem  to  be  due  neither  to  lack  of  oxygen  nor  to  the  excess 
of  carbon  dioxide  ;  rather  they  are  due  to  high  temperature 
and  to  the  presence  of  an  abnormal  amount  of  water  vapor, 
both  of  which  conditions  are  apt  to  prevail  in  crowded 
and  poorly  ventilated  rooms. 

Not  only  is  water  vapor  exhaled  from  the  lungs,  but 
there  is  constant  evaporation  of  moisture  from  the  pores 
of  the  skin,  and  in  this  process  much  heat  is  absorbed. 
Notwithstanding  the  extreme  changes  in  the  temperature 
of  the  air,  the  temperature  of  the  body  in  health  remains 


THE  ATMOSPHERE  141 

nearly  constant.  It  is  partly  by  variations  in  the  amount 
of  moisture  evaporating  from  the  skin  that  the  tempera- 
ture of  the  body  is  maintained  at  this  constant  value. 
If  an  abnormal  amount  of  water  vapor  is  present  in  the 
ah*,  the  evaporation  of  moisture  from  the  skin  takes 
place  very  slowly,  and  bodily  discomfort  follows.  More- 
over, when  the  air  is  perfectly  still,  that  portion  of  the 
air  in  contact  with  the  body  tends  to  become  saturated 
with  moisture,  and  evaporation  diminishes ;  hence  the 
relief  that  comes  from  keeping  the  air  in  motion,  as  with 
an  electric  fan. 

In  general,  a  moisture  content  of  about  70  per  cent  of 
that  required  for  saturation  is  most  conducive  to  comfort. 
The  volume  of  fresh  air  necessary  for  good  ventilation 
varies  greatly  with  conditions,  but  in  general  may  be  said 
to  be  about  30  cu.  ft.  per  minute  for  each  person. 

The  properties  of  air.  Inasmuch  as  air  is  composed 
principally  of  a  mixture  of  oxygen  and  nitrogen,  which 
elements  have  already  been  discussed,  its  properties  may 
be  inferred  largely  from  those  of  the  two  gases.  One  liter 
weighs  1.2928  g. 

Liquid  air.  We  have  seen  (p.  108)  that  air,  like  all  other 
gases,  can  be  liquefied  by  the  combined  effect  of  pressure 
and  low  temperature.  Liquid  air  is  essentially  a  mixture 
of  liquid  nitrogen  (boiling  point,  — 195.7°)  and  liquid 
oxygen  (boiling  point,  —  182.9°) ;  hence  if  liquid  air  is 
allowed  to  evaporate,  the  nitrogen  tends  to  vaporize  first. 
Advantage  is  taken  of  this  difference  in  boiling  points  to 
separate  the  oxygen  and  nitrogen  from  each  other,  and 
the  method  serves  as  a  commercial  one  for  obtaining  the 
two  gases.  Liquid  air  is  also  employed  when  very  low 
temperatures  are  desired.  It  may  be  kept  for  several 
hours  by  storing  it  in  Dewar  flasks  (Fig.  42). 


142    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  extremely  low  temperature  of  liquid  air  may  be 
inferred  from  the  fact  that  mercury  when  cooled  by  it  is 
frozen  to  a  mass  so  hard  that  it  may  be  used  for  driving 
nails. 

EXERCISES 

1.  When  oxygen  and  nitrogen  are  mixed  in  the  proportion  in 
which  they  exist  in  the  atmosphere,  heat  is  neither  evolved  nor  ab- 
sorbed by  the  process.    What  important  point  does  this  suggest  ? 

2.  What   essential   constituent   of  the  air  is  found  in  larger 
amount  in  manufacturing  districts  than  in  the  open  country  ? 

3.  Can  you  suggest  any  reason  why  the  growth  of  clover  in  a 
field  improves  the  soil  ? 

4.  When  ice  is  placed  in  a  vessel  containing  liquid  air,  the  latter 
boils  violently.    Explain. 

5.  Does  an  electric  fan  lower  the  temperature  of  a  room?   Why 
is  relief  obtained  by  its  use? 

6.  Would  combustion  be  more  intense  in  liquid  than  in  gaseous 
air? 

7.  Taking  the  volumes  of  the  oxygen  and  nitrogen  in  100  vol- 
umes of  air  as  21  and  78  respectively,  calculate  the  percentages  of 
these  elements  present  by  weight. 

8.  Assuming  that  dry  wood  contains  40  per  cent  carbon,  all  of 
which  originally  came  from  carbon  dioxide  in  the  air,  what  weight 
of  CO2  would  have  to  be  absorbed  by  a  plant  to  make  500  g.  of  wood  ? 

9.  What  weight  of  copper  would  be  required  to  combine  with 
the  oxygen  in  5  1.  of  dry  air? 

10.  A  tube  containing  calcium    chloride    was    found  to  weigh 
30.1293  g.    A  volume  of  air  which  weighed  15.2134  g.  was  passed 
through  the  tube,  after  which  the  weight  of  the  tube  was  found  to 
be  30.3405  g.    What  was  the  percentage  of  moisture  present  in  the 
air? 

11.  10 1.  of  air  measured  at  20°  and  740  mm.  passed  through  lime- 
water  caused  the  precipitation  of  0.0102  g.   of  CaCOs.     Find  the 
number  of  volumes  of  carbon  dioxide  in  10,000  volumes  of  the  air. 


CHAPTER  XIII 
SOLUTIONS;    THE  IONIZATION  THEORY 

Introduction.  In  Chapter  II  a  distinction  was  made  be- 
tween a  mixture  and  a  compound.  In  a  typical  mixture 
particles  of  different  properties  may  be  distinguished,  so 
it  is  not  of  perfectly  uniform  character.  In  a  compound 
every  smallest  portion  is  identical  in  composition  with 
every  other  portion. 

Intermediate  between  these  is  a  great  class  of  bodies 
called  solutions,  the  most  familiar  types  of  which  are  solu- 
tions of  solids  in  liquids.  They  differ  most  noticeably 
from  mixtures  in  that  they  are  of  perfectly  even  character 
throughout,  which  fact  is  usually  expressed  by  saying  that 
they  are  homogeneous.  They  differ  from  definite  chemical 
compounds  in  that  their  composition  can  be  varied  between 
wide  limits.  A  solution  may  therefore  be  defined  as  a  body 
of  homogeneous  character  whose  composition  may  be  varied 
continuously  between  certain  limits.  This  definition  makes 
no  restrictions  as  to  the  physical  state  of  the  solution  or 
of  its  constituents.  It  includes  any  combination,  such  as 
gases  in  gases  or  in  liquids,  solids  in  liquids  or  in  solids. 

In  all  solutions  we  are  dealing  with  two  constituents: 
(1)  the  medium  in  which  the  second  body  dissolves, 
which  is  known  as  the  solvent,  and  (2)  the  body  which 
dissolves  in  the  solvent,  known  as  the  solute ;  thus,  in  a 
solution  of  sugar  in  water,  the  water  is  the  solvent  and 
the  sugar  is  the  solute.  The  most  familiar  types  of  solu- 
tions are  those  of  gases  in  liquids  and  solids  in  liquids. 
143 


144    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

SOLUTION  OF  GASES  IN  LIQUIDS 

We  have  seen  that  oxygen,  hydrogen,  and  nitrogen  are 
slightly  soluble  in  water.  Accurate  study  has  led  to  the  con- 
clusion that  all  gases  are  soluble  to  some  extent,  not  only 
in  water  but  in  many  other  liquids.  The  amount  of  a  gas 
which  will  dissolve  in  a  liquid  depends  upon  a  number 
of  conditions,  and  these  can  best  be  understood  by  suppos- 
ing a  vessel  B  (Fig.  63)  to  be  filled  with  the  gas  and 
inverted  over  the  liquid  A.  Under  these  circumstances 
the  gas  cannot  escape  or  become  mixed  with  another  gas. 

Factors  affecting  the  solubility  of  gases.  A  number  of 
factors  affect  the  solubility  of  a  gas  in  a  liquid. 

1.  Nature  of  the  gas.  Other  conditions  being  equal,  each 
gas  has  its  own  peculiar  solubility,  just  as  it  has  its  own 
special  taste  or  odor.  The  solubility  of  gases  varies  between 
wide  limits,  as  will  be  seen  from  the  following  table, 
which  gives  the  volumes  of  some  well-known  gases  that 
will  dissolve  in  1  1.  of  water. 

SOLUBILITY  OF  GASES  IN  ONE  LITER  OF  WATER 


NAME  OF  GAS 

VOLUME  ABSORBED  AT  0°  AND 
I'XDER  700  MM.  PRESSURE 

Ammonia  _<.  *  , 
Hydrogen  chloride  
Sulfur  dioxide  .  .... 

1298.91. 
606.01. 
79  791 

4371 

Carbon  dioxide  
Oxygen  

1.7131. 
0.04961. 

Nitrogen  
Hydrogen  .-'...  .  . 

0.02331. 
0.02141. 

It  will  be  noted  that  some  gases,  such  as  ammonia,  are 
very  soluble.  This  is  due  to  the  fact  that  the  gas  in  part 
combines  with  the  water  to  form  a  compound. 


SOLUTIONS;  THE  KXNTZATION  THEORY     145 


2.  Nature  of  the   liquid.     The    character   of  the   liquid 
has  much  influence  upon  the   solubility  of   a  gas.    Each 
liquid,  such  as  water,  alcohol,  or  ether,  has  its  own  pecu- 
liar solvent  power.    From  the  solubility  of  a  gas  in  water 
no  prediction  can  be  made  as  to  its  solubility  in  other 
liquids. 

3.  Effect  of  pressure;  the  law  of  Henry.    Increase  of  pres- 
sure always  increases  the  weight  of  gas  going  inte  solution, 
the  increase  being  proportional  to  the  pressure.     This  is 
known  as  the  law  of  Henri/, 

having  been  formulated  by 
him  in  1803.  If  1  g.  of  a 
gas  dissolves  in  100  cc.  of 
water  at  atmospheric  pres- 
sure, 2  g.  will  dissolve  under 
2  atmospheres,  provided  the 
temperature  remains  con- 
stant. Under  high  pres- 
sure large  quantities  of  a 
gas  may  be  dissolved  in  a 
liquid.  In  such  solutions,  when  the  pressure  is  removed, 
that  fraction  of  the  gas  escapes  that  was  held  in  solution 
by  the  increased  pressure. 

4.  Influence  of  temperature.     In   general,  the   lower  the 
temperature  of  the  liquid,  the  larger  the  quantity  of  gas 
which  it  can  dissolve.    Thus,  1  1.  of  water  at  0°  will  dis- 
solve 0.0496  1.  of  oxygen  ;  at  50°,  0.01837  1. ;  at  100°,  none 
at  all.    While  most  gases  can  be  driven  from  a  liquid  by 
boiling  the  solution,  some  cannot.    For  example,  it  is  not 
possible  to  expel  hydrogen  chloride  completely  from  its 
solution  by  boiling. 


FIG.  63.   The  solubility  of  a  gas 
inclosed  over  water 


146    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

SOLUTIONS  OF  SOLIDS  IN  LIQUIDS 

A  solid  dissolved  in  a  liquid  is  by  far  the  most  familiar 
type  of  solution.  In  the  following  discussion  it  should  be 
remembered  that  we  are  dealing  with  true  solutions  only. 
Thus,  it  is  sometimes  said  that  zinc  dissolves  in  hydro- 
chloric acid.  In  this  case,  however,  the  solution  is  preceded 
by  an  undoubted  chemical  reaction  whereby  the  zinc  is 
converted  into  zinc  chloride,  and  it  is  this  compound  which 
is  obtained  when  the  solution  is  evaporated  to  dryness. 
With  solutions  such  .as  we  are  now  considering,  evapora- 
tion leaves  the  solute  in  its  original  chemical  condition. 

Factors  affecting  the  solubility  of  a  solid.  The  solubility 
of  a  solid  in  a  liquid  depends  upon  several  factors. 

1.  Nature  of  the  solid.    Other  conditions  being  the  same, 
solids  vary  greatly  in  their  solubility  in  liquids  (see  table 
on  opposite  page). 

No  solids  are  absolutely  insoluble,  but  the  amount  dis- 
solved may  be  so  small  as  to  be  of  no  significance  for  most 
purposes.  Thus,  barium  sulfate,  one  of  the  most  insoluble 
of  common  substances,  dissolves  in  water  to  the  extent  of 
1  part  in  400,000. 

2.  Nature  of  the  solvent.  Liquids  vary  much  in  their  power 
to  dissolve  solids.   Some  are  said  to  be  good  solvents,  since 
they  dissolve  a  great  variety  of  substances  and  considerable 
quantities  of  them.    Others  have  small  solvent  power,  dis- 
solving few  substances,  and  those  to  a  slight  extent  only. 
Broadly  speaking,  water  is  the  most  general  solvent,  and 
alcohol  is  perhaps  second  in  solvent  power. 

3.  Temperature.     The  weight  of  a  solid  which  a  given 
liquid  can  dissolve  varies  with  the  temperature.    Usually 
it  increases  rapidly  as  the  temperature  rises,  so  that  the 
boiling  liquid   dissolves  several  times  the  weight  which 


SOLUTIONS;  THE  IONIZATION  THEORY     147 


the  cold  liquid  will  dissolve.  In  some  instances,  as  in  the 
case  of  sodium  chloride  (common  salt)  dissolved  in  water, 
the  temperature  has  little  influence  upon  the  solubility, 
and  a  few  solids  such  as  calcium  hydroxide  (slaked  lime) 
are  more  soluble  in  cold  water  than  in  hot.  The  following 
examples  will  serve  as  illustrations: 

TABLE  OF  SOLUBILITY  OF  SOLIDS 


WEIGHT  DISSOLVED  BY  100  cc.  OF  WATER  AT 

0° 

20° 

100° 

Calcium  chloride 
Sodium  chloride 
Potassium  nitrate    . 
Copper  sulf  ate    .     . 
Calcium  sulfate  .     . 
Calcium  hydroxide  . 

CaCl2 
NaCl 
KNO3 
CuS04 
CaS04 
Ca(OH)2 

59.50  g.' 
35.70g. 
13.30  g. 
14.30  g. 
0.759  g. 
0.185  g. 

*  74.5  g. 
36.0  g. 
31.  6  g. 
21.7  g. 
0.203  g. 
0.165g. 

159.0g. 
39.80  g. 
246.0  g. 
75.4  g. 
0.162g. 
0.077  g. 

Molar  solutions.  In  stating  the  concentration  of  a  solu- 
tion we  may  obviously  make  use  of  the  percentage  sys- 
tem. It  is  often  desirable  to  state  the  number  of  formula 
weights  or  molecular  weights  (measured  in  grams)  which  a 
given  volume  of  the  solution  contains. 
When  as  many  grams  of  a  substance  as 
there  are  units  in  its  molecular  weight 
is  dissolved  so  as  to  make  a  liter  of  so- 
lution, it  is  said  to  be  a  molar  or 
gram-molecular  solution.  Thus,  a  molar 
solution  of  sodium  chloride  (NaCl) 
contains  23.00  +  35.46,  or  58.46  g.,  of 
the  solid  in  1  1.  of  the  solution. 

Saturated  solutions.  When  a  lump  of  sugar  is  placed 
in  a  small  beaker  and  covered  with  water,  as  represented 
in  Fig.  64,  it  gradually  diminishes  in  size  and  passes  into 


148    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


solution,  particles  leaving  it  and  diffusing  through  the 
solvent.  If  there  is  enough  sugar,  and  a  long  enough 
time  elapses,  the  concentration  of  the  sugar  in  the  solu- 
tion reaches  a  definite  limiting  value,  and  we  say  that  the 
sugar  ceases  to  dissolve  and  that  the  solution  is  saturated. 
There  is  good  reason  for  thinking  that  particles  continue 
to  leave  the  lump,  but  that  an 
equilibrium  has  been  reached, 
the  rate  of  departure  of  the  par- 
ticles being  equal  to  the  rate  of 
the  return  of  others  from  the  so- 
lution. A  saturated  solution  may 
therefore  be  defined  as  one  which 
is  in  equilibrium  with  the  undis- 
solved  solute. 

Separation  of  solids  from  their 
solutions.  When  any  given  solid 
separates  from  its  solution,  it  is 
usually  deposited  in  crystalline 
form,  each  solid  having  its  own 
definite  form.  This  process  is 
often  employed  to  purify  solids, 
such  solvents  being  selected  as 
will  retain  the  impurities  in 
solution.  The  separation  of  a  solid  from  its  solution  is 
usually  brought  about  in  one  of  two  ways:  (1)  The 
solvent  may  be  saturated  with  the  solid  at  room  tem- 
perature and  the  resulting  solution  set  aside.  The  solvent 
gradually  evaporates  and  the  solid  is  slowly  deposited. 
(2)  Most  solids  are  much  more  soluble  at  high  tempera- 
'tures  than  at  low ;  consequently,  if  a  solution  of  such  a 
solid  is  saturated  at  a  high  temperature  and  set  aside 
to  cool,  then,  in  most  cases  (see  next  paragraph),  all  of  the 


FIG.  65.  The  rapid  growth  of 
a  crystal  suspended  in  a  super- 
saturated solution 


SOLUTIONS;  THE  IONIZATION  THEORY     149 


solid  in  excess  of  the   quantity  required  to  saturate  the 
solution  at  the  lower  temperature  crystallizes  out. 

Supersaturated  solutions.  When  a  solution,  saturated 
at  a  given  temperature,  is  allowed  to  cool,  it  sometimes 
happens  that  no  solid  crystallizes  out,  although  the  solid 
may  be  much  less  soluble  in  the  cold  than  in  the  hot 
liquid.  This  is  very  likely  to  occur  when  the  solution 
is  rather  viscous,  like  sirup,  and  is  not  disturbed  in 
any  way.  Such  a  solution  is  said  to  be  supersaturated. 
That  this  con- 
dition is  un- 
stable can  be 
shown  by  add- 
ing a  crystal 
of  the  solid 
to  the  solu- 
tion. All  of 
the  solid  in 
excess  of  the 
quantity  re- 
quired to  saturate  the  solution  at  this  temperature  will  at 
once  crystallize  out  (Fig.  65).  Supersaturation  may  also 
be  overcome  in  many  cases  by  vigorously  shaking  or  stir- 
ring the  solution. 

Solubility  curves.  As  a  rule,  the  solubility  of  a  solute  does 
not  vary  with  temperature  in  any  regular  manner,  so  we  can 
best  represent  the  facts  in  a  given  case  by  a  curve,  plotting  the 
temperature  as  the  abscissa  (0  T)  and  the  number  of  grams  dis- 
solved in  1 1.  as  the  ordinate  (0  C).  The  diagram  (Fig.  66) 
shows  a  few  typical  curves.  It  will  be  seen  that  some  sub- 
stances are  very  soluble,  while  others  are  not.  The  solubility 
of  some  increases  rapidly  with  the  increase  in  temperature, 
while  with  others  the  increase  is  small. 


80°    90°    WO°T 


FIG.  66.  Some  typical  solubility  curves 


150    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Characteristic  properties  of  solutions.  A  few  general 
statements  may  be  made  in  reference  to  some  character- 
istic properties  of  solutions. 

1.  Distribution  of  the  solid  in  the  liquid.    A  solid,  when 
dissolved,  tends  to  distribute  itself  uniformly  through  the 
liquid,  so  that  every  part  of  the  solution  has  the  same 
concentration.     The   process  goes  on  very  slowly  unless 
hastened  by  stirring  or  shaking  the  solution.     If  a  few 
crystals  of  a  highly  colored  substance,  such  as  potassium 
permanganate,  are  placed  in  the  bottom  of  a  tall  vessel 
full  of  water,  it  will  take  weeks  for  the  solution  to  become 
uniformly  colored. 

2.  Boiling  point  of  solutions.   The  boiling  point  of  a  liquid 
is  raised  by  dissolving  a  solid  in  the  liquid.    In  general, 
the   extent  of  this  rise  is  proportional  to  the  molecular 
concentration  of  the  solution ;  that  is,  to  the  number  of 
gram-molecular  weights  of  the   substance  dissolved  in  a 
definite  weight  of  the  solvent.    It  appears,  therefore,  that 
it  is  not  the  character  of  the  molecules,  but  their  number 
alone,  which  affects  the  change  in  the  boiling  point. 

3.  Freezing  point  of  solutions.     The  freezing  point  of  a 
liquid  is  lowered  by  the  presence  of  a  substance  dissolved 
in  it.     The  lowering   of  the  freezing  point  obeys  a  law 
similar  to  the  one  which  holds  for  the  raising  of  the  boil- 
ing point,  the  extent  of  the  lowering  being  proportional  to 
the  molecular  concentration  of  the  solution. 

Electrolysis.  Pure  water  does  not  appreciably  conduct 
the  electric  current.  Moreover,  if  certain  compounds  such 
as  sugar  are  dissolved  in  the  water,  the  solution  is  also 
a  nonconductor.  If,  however,  certain  other  compounds 
such  as  sodium  chloride  or  sulfuric  acid  are  dissolved  in 
the  water,  the  resulting  solutions  are  found  to  be  good 
conductors  and  are  called  electrolytes.  When  the  current 


SOLUTIONS;  THE  IONIZATIOK  THEORY     151 

passes  through  an  electrolyte,  some  chemical  change  always 
takes  place.    This  change  is  called  electrolysis. 

The  general  method  used  in  the  electrolysis  of  a  solution 
is  illustrated  in  Fig.  67.  Two  plates  or  rods  (A  and  i?), 
made  of  suitable  material,  are  connected  with  the  wires  from 
a  battery  (or  dynamo)  C  and  dipped  into  the  electrolyte,  as 
shown  in  the  figure.  These  plates  or  rods  are  called  elec- 
trodes. The  electrode  B  connected  with  the  negative  pole 
of  the  battery  is  the  negative  electrode,  or  cathode,  while 
that  connected  with  the  positive 
pole  A  is  the  positive  electrode, 
or  anode. 

The  theory  of  ionization.  Just 
why  solutions  of  certain  com- 
pounds, such  as  sodium  chloride, 

ponduct  the  electric  current,  while     FlG:  6J"     DifgranJ 

method  of  electrolysis  of  a 
solutions    of    other    compounds,  solution 

such  as  sugar,  do  not,  is  not  defi- 
nitely known.  It  is  a  significant  fact  that  all  those  com- 
pounds whose  solutions  are  electrolytes  affect  the  boiling 
points  and  freezing  points  of  the  solvents  abnormally ;  that 
is,  the  boiling  points  and  the  freezing  points  are  changed 
more  than  we  should  expect.  The  solids  apparently  act 
as  though  their  molecules  when  dissolved  in  water  broke 
up  into  two  or  more  parts,  so  that  the  effect  is  the  same 
as  if  the  number  of  molecules  had  been  increased.  These 
facts,  taken  together  with  the  facts  discovered  in  connec- 
tion with  electrolysis,  are  best  explained  by  a  theory  first 
proposed  by  the  Swedish  chemist  Arrhenius  (Fig.  68),  and 
known  as  the  theory  of  ionization.  The  fundamental  ideas 
in  this  theory  are  as  follows: 

1.  Formation  of  ions.   The  molecules  of  many  compounds, 
when  dissolved  in  water,  fall  apart,  or  dissociate,  into  two 


152    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


or  more  parts,  called  ions.  Thus,  sodium  nitrate  (NaNOg) 
dissociates  into  the  ions  Na  and  NO8;  sodium  chloride 
(NaCl),  into  the  ions  Na  and  Cl.  These  ions  move  about 
in  the  solution  independently  of  each  other,  like  independ- 
ent molecules,  and  for  this  reason  were  given  the  name  ion, 

which  means  "  wanderer." 

2.  The  electrical  charge  of 
ions.    An   ion   differs    from 
an  atom  or  molecule  in  that 
it  carries  a   large  electrical 
charge.    It  is  evident  that 
the  sodium  ion  must  differ 
in  some  important  way  from 
ordinary  sodium,  for  sodium 
ions,  formed  from  ordinary 
salt,  give  no  visible  evidence 
of  their  presence  in  water, 
whereas  metallic  sodium  at 
once  decomposes  the  water. 
The  electrical  charge,  there- 
fore,   greatly    modifies    the 
usual  chemical  properties  of 
the  element. 

3.  The  positive  charges  equal 


FIG.  68.  Svante  August  Arrhenius 
(1859-) 


A  Swedish  chemist,  who  suggested 
the  theory  of  ionization 

the  negative  charges.   The  ions 

formed  by  the  dissociation  of  any  molecule  are  of  two 
kinds :  one  is  charged  with  positive  electricity  and  the 
other  with  negative.  The  sum  of  all  the  positive  charges 
is  always  equal  to  the  sum  of  all  the  negative  charges, 
and  the  solution  as  a  whole  is  therefore  electrically 
neutral.  If  we  represent  ionization  by  the  usual  chemical 
equations,  with  the  electrical  charges  indicated  by  plus  (+) 
and  minus  (— )  signs  following  the  symbols,  the  ionization 


SOLUTIONS;  THE  IONIZATION  THEORY     153 

of  the  molecules  of  sodium  chloride  and  sodium  sulfate 
is  represented  thus : 


Na+  +  Cl- 
Na+,  N 


NaCl- 

Na2SO4- 

Those  ions  that  are  positively  charged  are  known  as  cations, 

while  those  that  are  negatively  charged  are  termed  among. 

4.  Not  all  compounds  ionize.   It  is  assumed  that  only  those 

compounds  ionize  whose  solutions  are  electrolytes.    Thus, 


FIG.  69.    Apparatus  for  determining  whether  or  not  a  solution  is  an 
electrical  conductor 

salt  ionizes  when  dissolved  in  water,  for  it  has  been  found 
that  the  resulting  solution  is  a  very  good  electrolyte.  Sugar, 
on  the  other  hand,  does  not  ionize,  for  its  solution  is  not 
a  conductor  of  the  electric  current. 

Fig.  69  illustrates  a  very  convenient  apparatus  for  deter- 
mining whether  a  solution  is  a  good  conductor.  The  solution 
is  placed  in  the  bottle  A  and  the  electrodes  are  dipped  into  it. 
Connection  with  the  lighting  circuit  is  made  by  the  cord  and 
plug  B.  If  the  solution  is  a  good  conductor,  the  current  will 
flow  through  the  lamp  C,  which  will  then  glow. 


154    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Extent  of  ionization.  Compounds  differ  greatly  among 
themselves  in  the  ease  with  which  they  ionize  in  solution. 
Some  compounds  ionize  readily,  others  only  slightly,  and 
others  not  at  all.  Again,  the  extent  of  ionization  varies 
with  the  solvent.  Thus,  the  gas  known  as  hydrogen  chlo- 
ride (HC1)  is  largely  ionized  when  dissolved  in  water, 
while  in  benzene  it  is  not  ionized  at  all.  Water  is  the 
best  ionizing  solvent.  Moreover,  dilution  aids  ionization. 
For  example,  in  a  concentrated  aqueous  solution  of  sodium 
chloride,  only  a  small  percentage  of  the  molecules  are 
ionized,  an  equilibrium  being  reached  between  the  mole- 
cules decomposing  into  ions  and  those  forming  again  from 
the  ions,  as  expressed  in  the  following  equation  : 


If  the  solution  is  diluted  by  the  addition  of  more  water, 
the  percentage  of  molecules  undergoing  ionization  increases, 
and  reaches  a  maximum  only  in  very  dilute  solutions. 

The  theory  of  ionization  and  the  properties  of  solutions. 
In  order  to  be  of  value  this  theory  must  be  in  accord 
with  the  chief  properties  of  solutions.  Let  us  now  see  if 
the  theory  is  in  harmony  with  certain  of  these  properties. 

The  theory  of  ionization  and  the  boiling  and  freezing  points 
of  solutions.  We  have  seen  that  the  boiling  point  of  a 
liquid  is  raised  in  proportion  to  the  number  of  molecules 
of  a  solute  dissolved  in  the  liquid.  It  has  been  found  that 
in  the  case  of  electrolytes  the  boiling  point  is  raised  more 
than  it  should  be  to  conform  to  this  law.  If  the  solute 
dissociates  into  ions,  the  reason  for  this  becomes  clear. 
Each  ion  has  the  same  effect  on  the  boiling  point  that  a 
molecule  has,  and  since  their  number  is  always  greater 
than  that  of  the  molecules  from  which  they  were  formed, 
the  effect  on  the  boiling  point  is  abnormally  great. 


SOLUTIONS;  THE  IONIZATION  THEORY     155 

In  a  similar  way  the  theory -furnishes  an  explanation  of 
the  abnormal  lowering  of  the  freezing  point  of  electrolytes. 

The  theory  of  ionization  and  electrolysis.  The  changes 
taking  place  during  electrolysis  harmonize  very  completely 
with  the  theory  of  ionization.  This  will  become  clear  from 
a  study  of  the  following  examples: 

1.  Electrolysis  of  sodium  chloride.     Fig.  70  represents  a 
vessel  in  which  the   electrolyte  is  a  solution  of  sodium 
chloride  (NaCl).    According  to  the  theory  of  ionization, 
the  molecules   of  sodium 
chloride  dissociate  into  the 
ions  Na+  and  Cl~  as  soon 
as  the  compound  dissolves 
in  water.    Since  the  cath- 
ode B  has  a  large  negative 
charge   derived   from  the        FlG-  70-  Diagram  illustrating  the 
n    j/u       TVT  +     •  theory     of     the     electrolysis     of 

battery  C,  the   Na+    ions  sodium  chloride  (NaC1) 

are   attracted    to   it.     On 

coming  in  contact  with  the  cathode,  they  give  up  their 
positive  charge  and  are  then  ordinary  sodium  atoms.  They 
immediately  decompose  the  water  according  to  the  equation 

2  Na  +  2  H20 >-  2  NaOH  +  H2 

In  a  similar  way  the  chlorine  ions  (Cl~)  are  attracted  to 
the  positively  charged  anode  A,  and  upon  giving  up  their 
charge  to  it  they  are  set  free  as  chlorine  atoms  and  may 
either  combine  with  each  other  to  form  molecules  of  chlorine 
gas,  or  may  attack  the  water  as  represented  in  the  equation 

4  Cl  +  2  H2O  — >  4  HC1  +  O2 

It  is  to  be  carefully  noted  that  the  current  does  not  bring 
about  the  decomposition  of  the  solute  into  ions,  and  that  it  can 
pass  through  the  solution  only  when  ions  are  already  present. 


156    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

2.  Electrolysis  of  water.  The  reason  for  the  addition  of 
sulfuric  acid  to  water  in  the  preparation  of  oxygen  and 
hydrogen  by  electrolysis  (p.  16)  can  now  be  made  clear. 
Water  itself  is  not  an  electrolyte  to  an  appreciable  extent, 
for  it  does  not  form  enough  ions  to  carry  a  current.  Sul- 
furic acid  (H2SO4)  dissolved  in  water  is  an  electrolyte,  and 
dissociates  into  the  ions  2H+  and  SO4~~,  each  SO4~~  ion 
having  two  negative  charges.  In  the  process  of  electrolysis 
of  the  solution  the  H+  ions  travel  to  the  cathode,  and  on 
being  discharged,  escape  as  hydrogen  gas.  The  SO4~  ~  ions, 
when  discharged  at  the  anode,  act  upon  the  water,  setting 
free  oxygen  and  once  more  forming  sulfuric  acid : 

2  S04  +  2  H2O  — »-  2  H2SO4  +  O2 

The  sulfuric  acid  can  again  ionize  and  the  process  repeat 
itself  as  long  as  any  water  is  left.  Hence  the  hydrogen 
and  oxygen  set  free  in  the  electrolysis  of  water  really 
come  directly  from  the  acid  but  indirectly  from  the  water. 

Properties  of  electrolytes  dependent  upon  the  ions  present. 
When  a  substance  capable  of  forming  ions  is  dissolved  in 
water,  the  properties  of  the  solution  will  depend  upon  two 
factors :  (1)  the  ions  formed  from  the  substance ;  (2)  the 
undissociated  molecules.  Since  the  ions  are  usually  more 
active  chemically  than  the  molecules,  most  of  the  chemical 
properties  of  an  electrolyte  are  due  to  the  ions  rather  than 
to  the  molecules. 

The  solutions  of  any  two  substances  which  give  the 
same  ion  will  have  certain  properties  in  common.  Thus, 
all  solutions  containing  the  copper  ion  Cu++  are  blue, 
unless  the  color  is  modified  by  the  presence  of  ions  or 
molecules  having  some  other  color;  aqueous  solutions  of 
all  chlorides  give  a  precipitate  with  silver  nitrate,  since  all 
such  solutions  contain  the  chlorine  ion  Cl~. 


SOLUTIONS;  THE  IONIZATION  THEORY     157 

Source  of  the  charges  upon  the  ions.  There  has  been  much 
speculation  to  account  for  the  electrical  charge  of  the  ions.  It 
appears  to  be  very  probable  that  the  atoms  of  the  elements  are 
not  homogeneous  bodies  but  organized  systems,  each  contain- 
ing its  own  number  of  smaller  bodies,  which  are  called  electrons. 
These  electrons  are  all  alike,  and  appear  to  be  about  -J-^-Q  the 
weight  of  a  hydrogen  atom.  They  can  be  separated  from 
ordinary  matter,  so  they  are  capable  of  existing  in  the  free 
state.  The  evidence  goes  to  show  that  they  are  really  negative 
electricity,  which  is  therefore  a  material  thing.  A  body  con- 
taining more  than  its  normal  number  of  electrons  is  said  to  be 
negatively  charged,  while  one  from  which  some  of  its  normal 
number  of  electrons  have  been  removed  is  said  to  be  posi- 
tively charged.  Electrical  energy  is  the  energy  of  innumerable 
electrons  in  very  rapid  motion. 

Applying  these  views  to  the  electrification  of  ions,  we  as- 
sume that  before  union  the  atoms  of  sodium  and  chlorine  have 
each  their  normal  number  of  electrons.  When  these  combine  to 
form  sodium  chloride,  we  have  no  knowledge  as  to  any  disturb- 
ance in  the  distribution  of  the  electrons  in  the  several  atoms. 
When  the  sodium  chloride  is  dissolved  in  water,  however,  it 
appears  that  this  distribution  tends  to  change.  The  sodium 
atom  loses  one  electron  and  the  chlorine  atom  gains  one.  The 
sodium  atom  is  now  positively  charged,  the  chlorine  negatively ; 
and  in  this  condition  they  can  part  from  each  other  to  form 
independent  ions.  Upon  recombination  the  original  condition 
of  the  two  atoms  is  restored,  and  the  molecule  is  electrically 
neutral.  The  ions  are  therefore  very  different  things  from  the 
atoms;  they  should  even  have  different  weights,  though  we 
cannot  verify  this  experimentally. 


1.  Distinguish  clearly  between  the  following  terms:  electrolysis, 
electrolyte,  ions,  solute,  solvent,  solution,  saturated  solution,  and  super- 
saturated solution. 

2.  Why  does  the  water  from  some  natural  springs  effervesce? 

3.  Why  does  not  the  water  of  the  ocean  freeze  ? 


158    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

4.  Why  does  shaking  or  stirring  make  a  solid  dissolve  more 
rapidly  in  -a  liquid  ? 

5.  Why  will  vegetables  cook  faster  when  boiled  in  strong  salt 
water  than  when  boiled  in  ordinary  water  ? 

6.  How  do  you  explain  the  foaming  of  soda  water? 

7.  Account  for  the  fact  that  sugar  sometimes  deposits   from 
sirups,  even  when  no  evaporation  has  taken  place. 

8.  100  g.  of  sodium  chloride  was  dissolved  in  sufficient  water  to 
give  1500  cc.  of  solution.    What  was  the  molar  concentration  ? 

9.  200  g.  of  potassium  nitrate  was  dissolved  in  120  cc.  of  boiling 
water  and  the  solution  cooled  to  20°.    What  weight  of  potassium 
nitrate  separated? 

10.  A  saturated  solution  of  each  of  the  compounds  potassium 
nitrate,  sodium  chloride,  and  calcium  hydroxide  was  prepared  by 
heating  the  solids  with  1000  cc.  of  water.    How  would  the  lowering 
of  the  temperature  affect  each  solution? 

11.  10  g.  of  common  salt  was  dissolved  in  water  and  the  solution 
evaporated  to  dryness ;  what  weight  of  solid  was  left  ?    10  g.  of  zinc 
was    dissolved    in    sulfuric    acid    and    the    solution  evaporated  to 
dryness ;  what  weight  of  solid  was  left  ? 

12.  What  is  the  action  of  sodium  upon  water?    How  do  you 
account  for  the  fact  that  the  sodium  ions  present  in  a  solution  of 
sodium  chloride  do  not  decompose  the  water? 

13.  The  SO4  ion  always  bears  a  double  charge  of  electricity  as 
represented  by  the  expression  SO4 — .    What  action  occurs  when  this 
ion  gives  up  its  charge  of  electricity,  as  in  the  electrolysis  of  dilute 
sulfuric  acid? 

14.  What  changes  should  you  expect  to  take  place  if  a  current  of 
electricity  were  passed  through  an  aqueous  solution  of  copper  sul- 
fate  (CuSO4)  ?   Illustrate  by  diagram  similar  to  Fig.  70. 

15.  Contrast  the  effects  in  composition  produced  by  adding  water 
to  concentrated  solutions  of  sugar  and  salt  respectively. 

16.  How  do  you  account  for  the  fact  that  solutions  of  certain 
compounds  often  change  in  color  upon  dilution  ? 


CHAPTER  XIV 
CHLORINE;   HYDROGEN  CHLORIDE;   HYDROCHLORIC  ACID 

History  and  occurrence  of  chlorine.  Scheele,  who  was  the 
first  to  obtain  oxygen  in  the  pure  state,  was  likewise  the 
first  to  isolate  the  element  chlorine  (1774).  He  obtained 
the  element  by  the  action  of  hydrochloric  acid  upon  man- 
ganese dioxide,  a  method  of  preparation  which  is  still  used. 
For  years  after  its  discovery  chlorine  was  regarded  as  a 
compound  of  hydrochloric  acid  with  oxygen,  but  in  1810 
the  English  chemist  Davy  proved  its  elementary  character. 

The  most  abundant  compound  of  chlorine  is  sodium 
chloride,  or  common  salt.  This  compound  is  found  in  sea 
water  and  constitutes  large  solid  deposits  in  various  parts 
of  the  world  (rock  salt).  Chlorine  also  occurs  in  nature 
in  combination  with  potassium,  magnesium,  calcium,  and, 
to  a  limited  extent,  with  some  of  the  other  metals. 

Preparation.  The  laboratory  method  of  preparation  is 
different  from  the  commercial  method. 

1.  Laboratory  methods.  Two  methods  are  in  common  use 
for  the  preparation  of  chlorine  in  the  laboratory.  In  the 
first  of  these  hydrochloric  acid  (that  is,  an  aqueous  solu- 
tion of  hydrogen  chloride  (HC1))  is  warmed  with  manga- 
nese dioxide.  The  latter  compound  is  a  black  solid  and 
has  the  formula  MnO2.  Manganese  tetrachloride  (MnCl4), 
is  first  formed,  but  it  is  unstable  and  breaks  down  into 
manganous  chloride  (MnCl2)  and  chlorine  (C10),  thus: 

MnO2  +  4  HC1 >•  MnCl4  +  2  H,O 

MnCl4 >- 

169 


160    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Laboratory  apparatus.  The  manganese  dioxide  and  the  hy- 
drochloric acid  are  brought  together  in  a  flask  A  (Fig.  71), 
and  a  gentle  heat  is  applied.  The  chlorine  set  free  passes 
out  through  B,  bubbles  through  a  little  water  in  C  (which 

removes  any  hydro- 
chloric acid  carried 
over  with  it),  and 
finally  through  some 
sulfuric  acid  in  7), 
which  dries  the 
gas.  Being  some- 
what soluble  in 
water,  it  is  col- 
lected in  cylinder 
E  by  displacement 
of  air,  the  color 
showing  when  the 


FIG.  71.   The  preparation  of  pure  chlorine 


cylinder  is  full. 


The  second  method,  often  used  when  a  limited  supply 
of  chlorine  is  desired,  consists  in  oxidizing  hydrogen  chlo- 
ride according  to  the  equation 

4HC1  +  O2 — )-2H20  +  2Cl2 

Free  oxygen  may  be  used,  but  the  yield  is  small.  The 
use  of  suitable  catalytic  agents  (such  as  copper  sulfate) 
increases  the  yield,  but  it  is  much  more  convenient  to 
substitute  an  oxidizing  agent  for  the  free  oxygen.  The 
best  oxidizing  agent  for  the  purpose  is  the  purple-black 
solid  having  the  formula  KMnO4  and  known  as  potassium 
permanganate.  The  complete  reaction  is  complex  and  will 
be  explained  more  fully  in  a  later  chapter ;  for  the  present 
it  is  only  necessary  to  remember  that  in  a  general  way 
the  chlorine  is  liberated  according  to  the  above  equation, 
the  oxygen  being  derived  from  the  potassium  permanga- 
nate. While  this  method  of  preparation  is  somewhat  more 


CHLORINE 


161 


expensive  than  the  first  method,  it  nevertheless  has  an 
advantage  in  that  no  heat  is  required  and  the  rate  at 
which  chlorine  is  formed  is  easily  regulated. 

Laboratory  apparatus.  To  obtain  chlorine  in  the  laboratory 
by  this  method  the  potassium  permanganate  is  placed  in  a 
flask  A  (Fig.  72),  and  a  mixture  of  equal  volumes  of  concen- 
trated hydrochloric  acid  and  water  is  added,  drop  by  drop, 
from  a  funnel  (B)  provided  with  a  stopcock  so  that  the  flow  of 
the  liquid  can  be  regulated.  (Such 
a  funnel  is  called  a  separatory  fun- 
nel^) The  reaction  takes  place  at 
once,  and  the  chlorine  evolved 
through  C  is  collected  by  displace- 
ment of  air.  If  a  high  degree  of 
purity  is  desired,  the  gas  should 
be  bubbled  through  sulfuric  acid. 


2.  Commercial  method.  It  will 
be  recalled  that  when  a  solution 
of  sodium  chloride  is  electro- 
lyzed,  chlorine  is  evolved  at  the 
anode,  while  sodium  hydroxide 
(NaOH)  is  formed  at  the  cathode 
and  remains  in  solution  (p.  155). 


FIG.  72.  Preparation  of  chlo- 
rine from  potassium  perman- 
ganate and  hydrochloric  acid 


All  of  the  chlorine  prepared  for  commercial  purposes  in 
the  United  States  is  obtained  in  this  way.  The  method 
has  the  advantage  that  sodium  chloride  is  cheap,  and  that 
the  sodium  hydroxide  formed  in  the  process  has  many  com- 
mercial uses.  The  chlorine  so  obtained  is  either  compressed 
in  strong  iron  cylinders  and  shipped  in  this  form,  or  it  is 
passed  into  slaked  lime  and  forms  the  solid  known  as  chloride 
of  lime,  or  bleaching  powder  (p.  427),  which  can  be  easily 
shipped  and  from  which  the  chlorine  can  be  obtained 
as  needed. 


162    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Properties  of  chlorine.  Chlorine  is  a  greenish-yellow  gas 
which  has  a  peculiar,  suffocating  odor  and  produces  a  very 
irritating  effect  upon  the  throat  and  lungs.  Even  when 
inhaled  in  small  quantities,  it  often  produces  all  the  symp- 
toms of  a  very  hard  cold,  and  in  larger  quantities  may  have 
serious  and  even  fatal  effects.  Inhaling  ether  or  ammonia 
gives  some  relief.  Chlorine  is  nearly  2.5  times  as  heavy  as 
air,  and  can  therefore  be  col- 
lected by  displacement  of  air. 
One  volume  of  water  under 
ordinary  conditions  dissolves 
about  3  volumes  of  chlorine. 

Chemical  conduct  of  chlo- 
rine. At  ordinary  tempera- 
tures chlorine  is  one  of  the 
most  active  of  all  elements. 
Its  activity  may  be  described 
under  a  number  of  distinct 
headings. 

1.  Action  on  elements  other 
than  hydrogen.  Many  of  the 
elements  combine  directly  with 
chlorine  even  at  ordinary  tem- 
peratures. Thus,  phosphorus  burns  in  a  current  of  the 
gas,  while  antimony  and  arsenic,  in  the  form  of  a  fine 
powder,  at  once  burst  into  flame  when  dropped  into  jars 
of  the  gas  (Fig.  73).  When  a  strip  of  copper  foil  heated 
to  redness  is  dropped  into  chlorine,  the  two  elements  unite 
with  incandescence.  Even  inactive  metals  like  gold  and 
platinum  are  tarnished  by  the  gas.  The  compounds  formed 
by  the  union  of  chlorine  with  another  element  are  called 
chlorides,  and  constitute  a  very  important  group.  Common 
salt,  for  example,  is  chloride  of  sodium  (NaCl). 


FIG.  73.    The   burning  of  pow- 
dered metals  in  chlorine  gas 


CHLORINE 


163 


2.  Action  on  hydrogen.    Chlorine  has  a  strong  affinity  for 
hydrogen,  uniting  with  it  to  form  gaseous  hydrogen  chloride 
(HC1).    A  jet  of  hydrogen  burning  in  the  air  continues 
to  burn  when  introduced  into  a  jar  of  chlorine,  giving  a 
somewhat  luminous  flame 

(Fig.  74).  A  mixture  of 
the  two  gases  explodes 
violently  when  a  spark  is 
passed  through  it  or  when 
it  is  exposed  to  bright  sun- 
light. In  the  latter  case  it  is 
the  light  and  not  the  heat 
which  starts  the  action. 

3.  Action    on   compounds 
of  hydrogen.    Not  only  will 
chlorine   combine    directly 
with  free  hydrogen,  but  it 
will  often  abstract  the  ele- 
ment from  its  compounds. 
Thus,    when    chlorine     is 

passed  into  a  solution  of  hydrogen  sulfide,  the  chlorine 
combines  with  the  hydrogen,  while  sulfur  is  set  free  in 
accordance  with  the  following  equation  : 


FIG.  74.    The  burning  of  a  jet 

of  hydrogen  in  an  atmosphere 

of  chlorine 


The  same  tendency  is  very  strikingly  seen  in  the  action 
of  chlorine  upon  turpentine.  The  latter  substance  is  largely 
made  up  of  compounds  which  have  the  composition  repre- 
sented by  the  formula  C1QH16.  When  a  strip  of  paper 
moistened  with  warm  turpentine  is  placed  in  a  jar  of 
chlorine,  the  hydrogen  of  the  turpentine  combines  with 
the  chlorine  to  form  hydrogen  chloride,  and  the  carbon  is 
set  free  h?.  the  form  of  a  black  solid. 


164    AN  ELEMENTAKY  STUDY  OF  CHEMISTliY 


4.  Action  upon  water.  The  reaction  between  chlorine  and 
water  is  of  great  importance,  since  the  use  of  chlorine  as 
a  bleaching  agent  is  based  primarily  upon  it.  When  the 
chlorine  is  passed  into  water  both  hydrochloric  acid  and 
hypochlorous  acid  (HC1O)  are  formed: 

C12  +  H2O  -3=±  HC1  +  HC10 

The  hypochlorous  acid,  however,  is  a  very  unstable  com- 
pound and  breaks  down  slowly  in  the  dark  but  rapidly 
in  the  sunlight,  as  follows: 

2HC10  — 


Oo 


The    effect   of   sunlight  in  increasing 
the  action  of  chlorine  upon  water  may  be 
shown  in  the  following  way :  If  a  long 
tube   of   rather   large  diameter   is  filled 
with  a  saturated  solution  of  chlorine  in 
water   and  inverted    in   a   vessel  of  the 
same  solution  (as  shown  in  Fig.  75),  and 
the  apparatus  is  placed  in  bright  sunlight, 
bubbles  of  gas  will  soon  be  seen  to  rise 
FIG.  75.    Decomposi-    through  the  solution  and  collect  in  the 
tion  of  water  by  chlo-    tube      ^n  examination  of  this  gas  will 
rine  in  the  sunlight       show  ^  .<.  ^  Qxygen 

The  decomposition  of  water  through  the  action  of  chlo- 
rine is  also  greatly  increased  in  the  presence  of  some  sub- 
stance which  combines  with  the  oxygen  as  fast  as  it  is  set 
free.  Consequently  a  solution  of  chlorine  in  water  is  a 
good  oxidizing  agent,  and,  indeed,  it  is  often  used  as  such. 

5.  Action  upon  color  substances ;  bleaching  action.  Chlorine 
possesses  a  powerful  bleaching  action.  Strips  of  highly 
colored  cloth,  when  moistened  with  water  and  placed  in 
jars  of  chlorine,  rapidly  lose  their  color.  The  presence 
of  water  is  essential  to  the  change,  as  may  be  shown  by 


CHLORINE 


165 


placing  strips  of  the  dry  cloth  in  chlorine  from  which  the 
moisture  has  been  removed  by  bubbling  it  through  sul- 
furic  acid.  Under  these  conditions  the  color  of  the  cloth 
remains  unchanged.  It  is  probable  that  the  bleaching  action 
of  chlorine  consists  first  in  its  reaction  with  water  to  form 
hypochlorous  acid  (p.  164).  This  acid  then  decomposes, 
the  resulting  oxygen  reacting  with  the  color  substance  of 
the  cloth  to  form  colorless 
compounds.  It  is  evident, 
therefore,  that  chlorine  will 
bleach  only  those  materials 
the  coloring  matters  of 
which  are  changed  by  its 
action  into  colorless  com- 
pounds. It  has  no  bleach- 
ing action  on  such  color 
substances  as  carbon,  and 
hence  does  not  affect 
printer's  ink  made  from 
carbon.  It  cannot  be  used 
for  bleaching  certain  sub- 
stances like  silk  and  straw, 
since  it  injures  the  fabric. 


FIG.  76.   Bleaching  colored  cloths  by 
moist  chlorine 


Certain  foods,  such  as  dried  fruits  and  the  lower  grades  of 
flour,  are  also  bleached.  The  bleaching  of  flour  was  formerly 
prohibited,  but  is  now  largely  practiced,  since  it  enables  the 
miller  to  obtain  a  larger  yield  of  white  flour. 

Fig.  76  illustrates  the  bleaching  action  of  chlorine.  Strips 
from  the  same  piece  of  cloth  are  suspended  in  three  jars,  of 
which  the  first  contains  air,  the  second  di*y  chlorine,  and  the 
third  moist  chlorine.  It  will  be  noted  that  the  dry  chlorine 
has  almost  no  bleaching  action,  while  the  moist  chlorine  has 
partly  removed  the  color. 


166    AN  ELEMENTARY  STUDY  OF  CHEMISTEY 


6.  Action  as  a  germicide.  Chlorine  has  marked  germicidal 
properties,  and  the  free  element,  as  well  as  the  compounds 
from  which  it  is  easily  liberated,  are  used  as  disinfectants. 

Uses  of  chlorine.  As  has 
been  stated  above,  chlorine  is 
an  excellent  germicide  and 
bleaching  agent,  and  large 
quantities  of  the  element 
are  used  for  these  purposes. 
The  various  kinds  of  fabrics 
woven  from  vegetable  fibers, 
such  as  flax  and  cotton,  are 
always  more  or  less  colored 
by  the  presence  of  natural 
coloring  matter.  Hence,  if  a 
white  fabric  is  desired,  bleach- 
ing is  necessary.  This  was 
formerly  accomplished  by  ex- 
posing the  fabric  to  the  action 
of  the  air  and  sunlight,  but 
many  days  were  required  for 
the  completion  of  the  process. 
The  same  results  are  now 
obtained  in  a  very  short  time 
by  the  use  of  chlorine. 

As  a  rule  chlorine  is  used 
commercially  in  the  form  of 
bleaching  powder.  The  chlorine  present  in  this  substance 
can  be  liberated  easily  and  utilized  as  desired.  Increasing 
amounts  of  the  free  element  are  being  used  in  manufactur- 
ing processes  and  as  an  agent  for  destroying  the  microor- 
ganisms in  city  water  supplies  (Fig.  77).  Many  of  the 
gas  bombs  used  in  the  great  war  contained  chlorine. 


FIG.  77.  Chlorine  stored  in  bombs, 
for  use  in  water  purification 


HYDKOGEN  CHLORIDE  167 

Nascent  state.  It  will  be  noticed  that  when  oxygen  is 
set  free  from  water  by  chlorine,  it  is  at  that  instant  able 
to  do  what  ordinary  oxygen  gas  cannot  do,  for  it  bleaches 
substances  which  would  remain  unchanged  in  dry  air  or 
pure  oxygen.  It  is  generally  true  that  the  activity  of  an 
element  is  greatest  at  the  instant  of  its  liberation  from  its 
compounds.  To  express  this  fact,  elements  at  the  instant 
of  liberation  are  said  to  be  in  the  nascent  state,  the  word 
nascent  being  derived  from  a  Latin  word  meaning  "  to  be 
born."  This  greater  activity  is  usually  accounted  for  by 
supposing  that  an  element  at  the  instant  of  liberation  from 
a  compound  is  in  the  form  of  atoms,  and  is  therefore 
more  reactive  than  after  the  atoms  have  combined  to  form 
molecules. 

Hydrogen  chloride  (HC1).  We  have  seen  that  hydrogen 
and  chlorine  combine  directly  (Fig.  74)  to  form  the  gase- 
ous compound  known  as  hydrogen  chloride.  In  the  labora- 
tory it  is  more  convenient  to  prepare  this  compound  in  a 
different  way,  as  described  in  the  following  paragraph. 

Laboratory  preparation.  Hydrogen  chloride  is  ordinarily 
prepared  in  the  laboratory  by  the  action  of  sulfuric  acid 
upon  sodium  chloride.  Sodium  sulfate  (Na2SO4)  and  hydro- 
gen chloride  are  formed  according  to  the  following  equation : 

2  NaCl  +  H2SO4  — >-  Na2SO4  +  2  HC1 

It  will  be  noted  that  in  this  reaction  the  sodium  of  the 
sodium  chloride  exchanges  places  with  the  hydrogen  of 
the  sulfuric  acid.  Such  a  reaction  is  called  a  double 
decomposition. 

To  prepare  the  gas,  the  dry  sodium  chloride  is  placed  in  a 
flask  A  (Fig.  78),  sulfuric  acid  is  added,  and  the  flask  gently 
warmed.  The  hydrogen  chloride  is  rapidly  given  off  and  can 
be  collected  by  displacement  of  air.  To  prepare  a  solution  of 


168    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


the  gas  the  end  of  the  delivery  tube  is  fixed  just  above  the 
level  of  some  water  in  the  cylinder  B.  The  gas  is  very  soluble 
and  is  absorbed  as  fast  as  it  escapes  from  the  tube.  The  sodium 
sulfate  formed  is  a  white  solid  and  remains  in  the  flask  A. 

Commercial  preparation.  The  laboratory  method  is  like- 
wise used  in  the  preparation  of  hydrogen  chloride  on  a  com- 
mercial scale.  The  compound  is  also  prepared  commercially 
by  heating  sodium  chloride  with  the  white  solid  compound 
known  as  sodium  hydrogen  sulfate  (NaHSO4)  (p.  207) : 


NaCl  +  NaHSO, 


Na2SO4 


In  either  case  the  hydrogen  chloride  liberated  is  passed 
into  water,  in  which  it  readily  dissolves,  the  solution  con- 
stituting the  hydrochloric  add  of  commerce.  When  the 

materials  are  pure  a  color- 
less solution  is  obtained. 
The  most  concentrated  so- 
lution has  a  density  of 
about  1.2  and  contains 
approximately  43  per  cent 
of  hydrogen  chloride.  The 
ordinary,  or  commercial, 
hydrochloric  acid  sold  on 
the  market  is  usually  col- 
ored yellow  by  impurities. 
This  impure  acid  is  often 
called  muriatic  acid. 

Properties  of  hydrogen 
chloride.  Hydrogen  chlo- 
ride, a  colorless  gas,  is  1.26 

times  as  heavy  as  air.  When  inhaled,  it  has  an  irritating 
and  suffocating  effect.  At  0°  it  is  condensed  to  the  liquid 
state  by  a  pressure  of  28  atmospheres.  The  resulting  liquid 


FIG.  78.    The  preparation  of  a  solu- 
tion of  hydrogen  chloride 


HYDROGEN  CHLORIDE 


169 


is  colorless,  boils  at  -  82.9°,  and  solidifies  at  -  113°.  This 
liquid  does  not  conduct  electricity,  has  no  action  upon 
metals,  and  in  general  is  very  inactive.  Hydrogen  chloride 
is  very  soluble  in  water,  1  volume  of  the  latter  under 
standard  conditions  dissolving  506  volumes  of  the  gas. 
The  density  of  its  aqueous  solutions  increases  with  the 
amount  of  gas  dissolved,  as  shown  in  the  following  table, 
which  gives  the  percentage  by  weight  of  hydrogen  chloride 
present  in  solutions  of  various  densities,  the  measurements 
being  taken  at  15°. 


PER  CENT 

OF  HC1 

DENSITY 

PER  CENT 
OF  HC1 

DENSITY 

PER  CENT 
OFHCl 

DENSITY 

6.69 

1.0284 

20.04 

1.1006 

35'.02 

1.1779 

10.17 

1.0507 

25.06 

1.1265 

40.09 

1.2013 

15.22 

1.0761 

30.00 

1.1526 

43.40 

1.2134 

The  extreme  solubility  of  hydrogen  chloride  in  water  may 
be  shown  as  follows  :  A  perfectly  dry  flask  A  (Fig.  79)  is  filled* 
with  hydrogen  chloride.  This  flask  is  connected,  by  means  of 
a  glass  tube,  with  a  similar  flask  B,  which  is  nearly  filled  with 
water,  as  shown  in  the  figure.  The  end  of  the  tube  opening 
into  flask  A  is  drawn  out  to  a  rather  fine  jet.  By  blowing  into 
the  tube  C,  a  few  drops  of  water  are  forced  into  A.  Some  of 
the  hydrogen  chloride  at  once  dissolves,  thus  diminishing  the 
pressure  inside  the  flask.  The  water  then  flows  continuously 
from  B  into  A,  until  virtually  all  the  hydrogen  chloride  is 
absorbed.  It  is  evident  that  the  connection  must  be  air-tight. 

Composition  of  hydrogen  chloride.  The  composition  of 
hydrogen  chloride  can  be  determined  by  the  electrolysis 
of  its  aqueous  solution.  When  hydrogen  chloride  is  elec- 
trolyzed,  the  hydrogen  of  the  compound  is  evolved  at  the 
cathode  and  the  chlorine  at  the  anode.  A  special  form  of 
apparatus  is  required,  in  order  to  avoid  the  difficulties 


170    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

arising  from  the  marked  solubility  of  the  chlorine  in 
water.  When  the  experiment  is  carried  out,  it  is  found 
that  the  volume  of  the  hydrogen  liberated  is  exactly 
equal  to  that  of  the  chlorine.  Conversely,  it  is  possible 
to  show  by  experiment  that  when  hydrogen  and  chlorine 
combine,  they  always  do  so  in  the  ratio  of  1  volume  of 
hydrogen  to  1  volume  of  chlo- 
rine ;  moreover,  the  product  is 
always  2  volumes  of  hydrogen 
chloride.  These  relations  may 
be  shown  graphically  in  the 
following  way : 


CH 


C12  —  +  HC1  HC1 


Since  chlorine  is  35.18  times 
as  heavy  as  hydrogen,  it  follows 
that  1  part  by  weight  of  hy- 
drogen combines  with  35.18 
EIG.  79.  Apparatus  to  show  parts  by  weight  of  chlorine  to 
the  extreme  solubility  of  hy-  form  gg<18  ^  b  wej  ht  of 

drogen  chloride  in  water  A 

hydrogen  chloride. 

Hydrochloric  acid.  Hydrogen  chloride  itself,  as  well  as 
its  solution  in  rionionizing  solvents,  has  but  little  chemical 
activity.  On  the  other  hand,  its  solution  in  water,  namely 
hydrochloric  add,  has  marked  chemical  properties  and  con- 
stitutes one  of  the  most  important  members  of  that  general 
class  of  compounds  known  as  acids  (p.  179).  It  has  a 
sour  taste  and  changes  the  colors  of  certain  organic  com- 
pounds. Thus,  blue  litmus  (a  colored  compound  obtained 
from  certain  lichens)  is  turned  red  when  a  little  hydro- 
chloric acid  is  added  to  it. 


HYDKOCHLORIC  ACID  171 

Hydrochloric  acid  acts  upon  many  of  the  metals,  such 
as  iron,  zinc,  and  sodium,  liberating  the  hydrogen  and 
forming  chlorides  of  the  metals: 

Fe  +  2  HC1  —  >•  FeCl2  +  H2 

Zn  +  2  HC1  -  *  ZnCl2  +  H2 

2  Na  +  2  HC1  —  >-  2  NaCl  +  H2 

It  also  acts  upon  the  hydroxides  of  the  metals,  such  as 
sodium  hydroxide  (NaOH),  as  shown  in  the  following 
equation  : 


EXERCISES 

1.  Consult  the  dictionary  for  the  meaning  of  the  word  chlorine. 

2.  Why  must  chlorine  water  be  kept  in  the  dark  ? 

3.  Calculate  the  weight  of  10  1.  of  chlorine.    What  weight  of 
the  compound  HC1  would  be  required  for  the  preparation  of  10  1.  of 
chlorine  ? 

4.  Why  not  collect  chlorine  over  water  as  hydrogen  and  oxygen 
were  collected? 

5.  What  volume  of  chlorine  can  be  prepared  from  100  kg.  of 
sodium  chloride? 

6.  10  1.  of  chlorine  will  combine  with  how  many  liters  of  hy- 
drogen ?    How  many  liters  of  hydrogen  chloride  will  be  formed  ? 

7.  Distinguish  definitely  between  hydrogen  chloride  and  hydro- 
chloric acid. 

8.  What  weight  of  common  salt  is  necessary  for  the  preparation 
of  100  1.  of  hydrogen  chloride? 

9.  In  the  preparation  of  hydrogen  chloride  (Fig.  78)  why  not 
have  the  exit  tube  dip  below  the  surface  of  the  water  in  B? 

10.  (a)  What  is  the  density  of  hydrochloric  acid  containing  30 
per  cent  of  hydrogen  chloride?  (See  table,  p.  169.)  (6)  What  would 
1  1.  of  this  acid  weigh  ?  (c)  What  weight  of  hydrogen  chloride 
would  it  contain?  (rf)  What  weight  of  sodium  chloride  would  be 
necessary  to  prepare  this  weight  of  hydrogen  chloride? 


CHAPTER  XV 
SODIUM;    SODIUM  HYDROXIDE 

Metals  and  nonmetals.  The  chemist  finds  it  convenient 
to  divide  the  elements  into  two  general  groups  known  as 
the  metals  and  the  nonmetals.  It  is  the  chemical  conduct  of 
an  element  that  determines  to  which  of  these  two  groups 
it  belongs.  This  distinction  will  be  discussed  in  a  later 
chapter.  For  the  present  it  is  only  necessary  for  us  to 
remember  that  all  the  metals  are  solids,  except  mercury, 
which  is  a  liquid ;  that  as  a  rule  they  are  good  conductors 
of  heat  and  electricity  and,  with  the  exception  of  gold 
and  copper,  they  have  a  silvery  luster.  Most  of  the  metals 
have  a  high  density ;  a  few,  such  as  aluminium  and  mag- 
nesium, however,  are  comparatively  light,  while  three, 
namely  lithium,  sodium,  and  potassium,  are  so  light  that 
they  will  float  on  water. 

The  elements  so  far  studied,  namely  oxygen,  hydrogen, 
nitrogen,  the  rare  elements  in  the  atmosphere,  and  chlo- 
rine, are  all  nonmetals.  Having  studied  a  number  of  non- 
metals  it  is  advisable  now  to  study  some  metal,  and  the 
one  known  as  sodium  best  serves  our  purposes. 

History  of  sodium.  The  isolation  of  sodium  dates  back 
to  the  year  1807.  At  that  time  the  compounds  now  known 
as  sodium  hydroxide  and  potassium  hydroxide  were  well 
known,  but  they  were  regarded  as  elementary  in  character. 
In  1807  Sir  Humphry  Davy  (Fig.  80),  while  studying 
the  effect  of  the  electric  current  upon  various  substances, 
172 


SODIUM;  SODIUM  HYDROXIDE 


173 


succeeded  in  decomposing  these  compounds  and  from  them 
obtained  both  sodium  and  potassium. 

Occurrence  of  sodium.  Because  of  its  great  activity, 
sodium  does  not  occur  in  nature  in  a  free  state.  The  most 
familiar  compound  of  the  element  found  in  nature  is  sodium 
chloride.  This  is  a  constitu- 
ent of  all  sea  waters  and 
mineral  waters  and  forms 
large  solid  deposits  in  vari- 
ous parts  of  the  world.  The 
element  also  occurs  as  a  con- 
stituent of  many  rocks,  and 
its  compounds  are  therefore 
present  in  the  soil  formed  by 
their  disintegration.  Other 
compounds  of  sodium  often 
found  in  nature  are  sodium 
nitrate  (known  commer- 
cially as  Chile  saltpeter), 
sodium  carbonate,  and  so- 
dium borate,  or  borax. 

Preparation    of     sodium. 


FIG.  80.    Sir  Humphry  Davy 
(1778-1829) 


A  distinguished  English  scientist  who 

invented  the  safety  lamp,  demonstrated 

the  nature  of  chlorine,  and  first  isolated 

the  elements  sodium  and  potassium 


For   many   years  the  most 

economical  method  known 

for  preparing  sodium  consisted  in  heating  sodium  carbon- 

ate (a  white  solid  having  the  formula  Na2CO8)  with  carbon  : 


NaC0 


Na  +  SCO 


At  present  it  is  prepared  by  the  electrolysis  of  its  hydrox- 
ide, a  white  solid  compound  having  the  formula  NaOH. 
It  is  necessary  to  carry  on  the  process  in  the  absence  of 
water,  because  the  metal  acts  upon  water  with  great  vigor 
(p.  39).  Since  it  is  impossible  to  obtain  the  element  by 


174    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


the  electrolysis  of  a  solution  of  the  hydroxide,  the  solid 
compound  is  heated  until  it  melts,  and  the  current  of  elec- 
tricity is  then  passed  through  the  resulting  liquid.  Sodium 
is  prepared  at  Niagara  Falls,  where  water  power  is  utilized 
for  generating  the  electric  current.  The  process  is  a  diffi- 
cult one  to  carry  out,  so  that 
sodium  is  not  a  cheap  metal. 

Commercial  preparation  (Castner's 
process).  At  Niagara  Falls  sodium 
is  prepared  by  the  electrolysis  of 
fused  sodium  hydroxide  by  a  process 
devised  by  Castner.  The  apparatus 
consists  of  a  cylindrical  iron  vessel 
A,  A  (Fig.  81),  through  the  bottom 
of  which  extends  an  iron  rod  B, 
which  serves  as  the  cathode.  The 
iron  anodes  C,  C,  several  in  number, 
are  suspended  around  the  cathode, 
but  are  kept  from  touching  it  by  a 
cylinder  of  iron  gauze  D,  which  is 
fastened  to  the  vessel  E.  The  lower 
part  of  the  vessel  A,  A  is  filled  with 
molten  sodium  hydroxide,  which,  on 
cooling,  holds  the  cathode  in  posi- 
tion. The  heat  generated  by  the  current  is  ordinarily  sufficient 
to  keep  the  hydroxide  in  the  upper  portion  of  the  vessel 
fused ;  however,  the  apparatus  is  supplied  with  a  row  of  gas 
burners  G,  G,  which  may  be  utilized  if  additional  heat  is 
required.  Sodium  and  hydrogen  are  liberated  at  the  cathode 
and,  rising  to  the  surface,  collect  in  vessel  E.  The  hydrogen 
escapes  by  lifting  the  cover  of  the  vessel,  while  the  sodium, 
protected  from  the  air  by  the  hydrogen,  is  skimmed  or  drawn 
off  from  time  to  time.  Oxygen  is  liberated  at  the  anodes  and 
escapes  through  the  opening  F  without  coming  in  contact  with 
either  the  sodium  or  the  hydrogen.  Sodium  has  also  been  pre- 
pared by  the  electrolysis  of  fused  sodium  chloride. 


FIG.  81.    A  Castner  cell  for 

the  electrolytic  production 

of  metallic  sodium 


SODIUM;  SODIUM  HYDKOXIDE  175 

Properties  and  uses.  Sodium  is  a  soft,  silver-white  metal, 
slightly  lighter  than  water.  It  melts  at  97.5°  and  boils  at  877°. 
It  is  very  active  chemically,  combining  readily  with  most 
of  the  nonmetallic  elements  such  as  oxygen  and  chlorine. 
It  liberates  hydrogen  from  the  common  acids  (p.  179)  and 
decomposes  water,  forming  sodium  hydroxide  and  hydrogen 

(p.  39) : 

2  Na  +  2  H2O  — h  2  NaOH  +  H2 

When  exposed  to  the  air,  sodium  rapidly  tarnishes,  owing 
to  its  combination  with  oxygen  and  carbon  dioxide.  On 
this  account  it  is  often  kept  immersed  in  kerosene,  since 
it  has  no  action  upon  this  liquid. 

Compounds  of  sodium.  Sodium  forms  many  useful  com- 
pounds. One  of  these,  sodium  hydroxide,  is  a  typical 
member  of  that  important  class  of  compounds  known  as 
bases,  and  it  is  desirable  for  us  to  study  its  properties  at 
this  time;  the  discussion  of  the  other  compounds  of  sodium 
may  well  be  deferred  to  a  later  chapter. 

Sodium  hydroxide  (caustic  soda)  (NaOH).  This  compound 
is  a  white  solid  and,  as  we  have  seen,  may  be  prepared 
by  the  action  of  sodium  upon  water.  It  is  only  necessary 
to  bring  the  sodium  and  water  in  contact  with  each  other. 
The  resulting  sodium  hydroxide  dissolves  in  the  water  as 
fast  as  it  is  formed,  and  may  be  recovered  by  the  evapo- 
ration of  the  water.  Since  sodium  is  a  comparatively 
expensive  metal,  it  is  evident  that  this  method  is  not  an 
economical  one.  Two  general  processes  are  now  used  for 
preparing  the  compound  on  a  commercial  scale. 

1.  Action  of  calcium  hydroxide  upon  sodium  carbonate.  This 
process  consists  in  treating  calcium  hydroxide  (slaked  lime) 
suspended  in  water  with  sodium  carbonate : 

Na2CO8  +  Ca(OH)2  — >•  CaCO8  +  2  NaOH 


176    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


The  calcium  carbonate  (CaCO3)  is  precipitated  as  a  white, 
insoluble  solid,  leaving  the  sodium  hydroxide  in  solution. 

While  this  is  an  old  process  it  still  remains  the  chief  one 
for  the  production  of  the  hydroxide.  Manufacturers  of  sodium 
carbonate  often  utilize  a  portion  of  their  product  in  the  prepa- 
ration of  the  hydroxide,  so  that  the  manufacture  of  these  two 
compounds  is  often  carried  out  in  the  same  plant. 

2.  Electrolytic  methods.  By  the  newer 
method  sodium  hydroxide  is  obtained 
by  the  electrolysis  of  sodium  chloride. 
The  products  of  the  electrolysis  are 
sodium  hydroxide,  hydrogen,  and  chlo- 
rine (p.  155).  The  chief  difficulty  in 
this  process  is  to  prevent  the  chlorine 
and  the  hydroxide  from  acting  upon 
each  other.  This  is  usually  done  by 
separating,  by  means  of  a  porous 
partition,  the  anode  and  cathode  com- 
partments of  the  cell  in  which  the 
electrolysis  is  effected.  A  number  of 
different  cells  have  been  devised  for 
carrying  out  the  process ;  at  present 
one  of  the  most  successful  of  these 
is  that  devised  by  Townsend  and 
known  by  his  name. 


FIG.  82.    A  Townsend 

cell  for  the  production 

of  sodium  hydroxide 


Commercial  preparation  by  the  Townsend  cell.  A  section  of  this 
cell  is  shown  in  Fig.  82.  The  anode  compartment  is  formed  by 
the  partition,  or  diaphragm,  A,  A,  a,  nonconducting  bottom  B, 
and  a  lid  C.  The  diaphragm  is  made  of  asbestos  cloth  painted 
over  with  a  mixture  of  iron  oxide  and  asbestos  fiber.  Through 
the  lid  C  extends  the  graphite  anode  D.  The  diaphragm  is  set 
firmly  against  the  perforated  iron  cathode  plate  E,  E,  which 
is  in  turn  held  in  place  by  the  iron  sides  F,  F,  the  space 


SODIUM;  SODIUM  HYDROXIDE  177 

between  the  plate  and  the  iron  sides  forming  the  cathode 
compartment.  The  anode  compartment  is  partly  filled  with 
saturated  salt  solution  G,  and  the  cathode  compartment  with 
kerosene  H.  Since  the  level  of  the  salt  solution  is  above  that 
of  the  kerosene,  the  solution  slowly  penetrates  the  diaphragm, 
and  some  of  the  salt,  coming  in  contact  with  the  cathode,  is 
changed  into  the  hydroxide.  The  resulting  solution  of  the 
chloride  and  hydroxide  enters  the  anode  compartment  and, 
being  heavier  than  the  kerosene,  sinks  to  the  bottom  and  is 
drawn  off  through  the  side  tubes  I.  The  chloride,  being  much 
less  soluble  than  the  hydroxide,  is  separated  by  partial  evapo- 
ration of  the  solution.  The  hydrogen  and  the  chlorine  that  are 
set  free  are  led  off  through  tubes,  and  the  chlorine  is  used  in 
the  preparation  of  bleaching  powder. 

Properties  of  sodium  hydroxide.  Sodium  hydroxide  is  a 
brittle  white  crystalline  substance  which  rapidly  absorbs 
water  and  carbon  dioxide  from  the  air.  As  the  name 
caustic  soda  indicates,  it  is  a  very  corrosive  substance,  hav- 
ing a  disintegrating  action  on  most  animal  and  vegetable 
tissues.  It  is  very  soluble  in  water,  and  the  resulting  solu- 
tion is  soapy  to  the  touch  and  has  a  bitter  taste.  The 
solution  turns  red  litmus  blue,  reversing  the  color  change 
produced  on  litmus  by  hydrochloric  acid.  Its  reaction  with 
hydrochloric  acid  has  been  described  (p.  171).  It  is  used 
in  a  great  many  chemical  industries,  its  chief  use  being  in 
the  manufacture  of  soap.  As  a  household  article  it  is  sold 
under  the  name  of  lye. 


1.  What  is  the  word  from  which  the  symbol  of  sodium  is  derived? 
(Consult  dictionary.) 

2.  Is  sodium  one  of  the  abundant  elements?    (Consult  Clarke's 
table,  p.  10.) 

3.  What  is  the  approximate  weight  of  the  sodium  in  your  body? 


178    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

4.  What  weight  of  sodium  hydroxide  is  necessary  for  the  prepa- 
ration of  1  kg.  of  sodium  ? 

5.  When  sodium  acts  upon  water,  what  volume  of  hydrogen  is 
liberated  for  each  gram  of  sodium  consumed? 

6.  How  many  pounds  of  sodium  chloride  are  necessary  for  the 
preparation  of  1  ton  of  caustic  soda  ? 

7.  Calculate  the  percentage  composition  of  sodium  hydroxide. 

8.  Write  the  equation  for  the  reaction  between  sodium  hydroxide 
and  hydrochloric  acid. 

9.  What  weight  of  sodium  hydroxide  can  be  obtained  from 
500  g.  of  sodium  carbonate? 

10.  How  could  you  tell  whether  a  given  solution  contains  sodium 
hydroxide  or  hydrochloric  acid,  if  you  knew  that  one  of  these  is 
present  ? 


CHAPTER  XVI 
ACIDS,  BASES,  AND  SALTS;    NEUTRALIZATION 

Introduction.  The  great  majority  of  the  compounds  to 
be  described  in  the  course  of  our  study  belong  to  one  of 
three  classes,  known  as  acids,  bases,  and  salts.  We  have 
already  studied  the  properties  of  hydrochloric  acid,  as  well 
as  those  of  the  well-known  base,  sodium  hydroxide.  More- 
over, sodium  chloride,  formed  by  the  action  of  hydrochloric 
acid  upon  sodium  hydroxide,  is  a  typical  salt.  Having 
learned  some  facts  about  a  representative  of  each  of  these 
classes,  we  shall  now  proceed  to  a  more  careful  study  of 
the  characteristics  of  each  class. 

The  common  acids.  It  will  be  recalled  that  hydrochloric 
acid  consists  of  a  solution  of  hydrogen  chloride  in  water. 
All  of  the  other  common  acids  used  so  largely  in  the 
industries  and  in  chemical  laboratories  are  aqueous  solu- 
tions of  definite  compounds.  Thus,  nitric  acid  is  a  solution 
of  the  liquid  known  as  hydrogen  nitrate,  the  formula  of 
which  is  HNOg ;  sulfuric  acid  is  a  solution  of  the  thick 
oily  liquid  called  hydrogen  sulfate,  whose  formula  is  H2SO4. 
For  most  purposes  it  is  not  necessary  to  make  a  distinc- 
tion between  the  name  of  the  compound  and  its  solution 
in  water,  and  both  are  frequently  called  acids. 

Characteristics  of  acids.  A  study  of  the  acids  reveals  the 
fact  that,  while  the  individual  members  of  the  class  may 
differ  from  each  other  in  many  of  their  properties,  they  all 
possess  certain  fundamental  characteristics.  These  charac- 
teristics are  as  follows:  (1)  All  compounds  forming  acids 
179 


180    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

in  solution  contain  hydrogen.  (2)  They  react  with  the 
hydroxides  of  metals,  in  which  reaction  the  hydrogen  of 
the  acid  combines  with  the  hydroxyl  group  (OH)  of  the 
base  to  form  water.  Thus,  hydrochloric  acid  and  sulfuric 
acid  react  with  sodium  hydroxide  as  indicated  in  the 
following  equations  : 

HC1  +  NaOH  -  >-  NaCl  +  H2O  (or  HOH) 
H2SO4  +  2  NaOH  —  *  NaaSO4  +  2  H2O 

(3)  As  a  general  rule,  acids  act  upon  certain  metals,  dis- 
solving the  metals.  In  this  reaction  the  metal  displaces  the 
hydrogen  of  the  acid.  Thus,  zinc  reajcts  with  hydrochloric 
acid  and  sulfuric  acid  as  shown  in  the  following  equations  : 

Zn  +  2  HC1  -  >•  ZnCl2  +  H2 


Zn  +  H2SO4  -  >•  ZnSO4 

(4)  The  solutions  have  a  sour  taste  and  change  the  color 
of  certain  compounds  known  as  indicators.  Thus,  blue 
litmus  turns  red  on  the  addition  of  an  acid.  (5)  The 
solutions  are  electrolytes,  and  when  a  current  of  electricity 
is  passed  through  them,  hydrogen  is  evolved  at  the  cathode. 
The  compounds  forming  acids  ionize  when  dissolved  in 
water;  thus,  HC1—  >H+,  Cl- 

HNO3  —  >-H+,  NO8- 

H2SQ4  —  >-H+,  H+,  SO4— 

It  will  be  noted  that  the  one  constituent  common  to  all 
acids  is  the  hydrogen  ion,  H+.  Accordingly,  the  properties 
which  all  of  these  acids  have  in  common,  such  as  their 
action  on  metals  and  their  hydroxides,  must  be  due  to 
the  hydrogen  ions.  From  the  standpoint  of  the  ionization 
theory,  therefore,  we  may  define  an  acid  as  a  solution 
containing  hydrogen  ions. 


ACIDS,  BASES,  AND  SALTS  181 

Undissociated  acids.  When  compounds  whose  aqueous 
solutions  form  acids  are  perfectly  free  from  water,  they 
have  no  real  acid  properties.  Neither  do  they  have  acid 
properties  when  they  are  dissolved  in  liquids,  like  benzene, 
which  do  not  have  the  power  of  dissociating  them  into 
ions.  For  example,  a  benzene  solution  of  hydrogen  chloride 
has  no  action  on  litmus;  neither  does  it  dissolve  metals. 

The  common  bases.  In  addition  to  sodium  hydroxide,  the 
common  bases  include  potassium  hydroxide  (KOH)  and  cal- 
cium hydroxide,  Ca(OH)2.  These  are  white  solids  soluble 
in  water.  The  very  soluble  bases  with  most  pronounced 
basic  properties  are  often  called  alkalies.  While  these 
compounds  are  often  called  bases,  yet  pronounced  basic 
properties  are  exhibited  only  by  their  solutions  in  water. 

Characteristics  of  bases.  The  class  characteristics  of  the 
bases  are  as  follows:  (1)  The  compounds  contain  hydro- 
gen, oxygen,  and  a  metal.  (2)  They  react  with  acids,  the 
hydroxyl  group  (OH)  of  the  base  combining  with  the 
hydrogen  of  the  acid.  (3)  Their  solutions  reverse  the  color 
change  produced  in  indicators  by  acids;  thus,  they  turn 
red  litmus  blue.  (4)  Their  solutions  are  conductors  of 
electricity,  and  when  these  solutions  are  electrolyzed  the 
metal  is  set  free  at  the  cathode. 

When  dissolved  in  water,  the  molecules  of  the  base 
dissociate  into  two  kinds  of  ions.  One  of  these  is  always 
composed  of  the  group  (OH),  and  is  the  anion.  It  is  called 
the  hydroxyl  ion.  The  remainder  of  the  molecule,  which 
usually  consists  of  a  single  atom,  is  the  cation ;  thus, 

NaOH — >-Na+,  OH~ 

Since  all  bases  produce  hydroxyl  anions,  while  the 
cations  of  each  are  different,  the  properties  which  all  bases 
have  in  common  when  in  solution  must  be  due  to  the 


182    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

hydroxyl  ions.  From  the  standpoint  of  the  ionization  theory, 
therefore,  we  may  define  a  base  as  a  solution  containing 
hydroxyl  ions. 

Neutralization.  When  an  acid  and  a  base  are  brought 
together  in  solution  in  the  proper  proportion,  the  charac- 
teristic properties  of  each  disappear.  It  is  evident,  there- 
fore, that  the  acid  and  base  must  have  reacted  with 
each  other  to  form  one  or  more  new  compounds  which 
are  neither  acid  nor  basic  in  character.  What  really  hap- 
pens is  that  the  hydrogen  of  the  acid  and  the  hydroxyl 
group  (OH)  of  the  base  combine  to  form  water,  while 
the  remaining  constituents  of  the  acid  and  base  unite  to 
form  a  compound  which  belongs  to  the  class  of  com- 
pounds known  as  salts.  These  facts  are  shown  by  the 
following  typical  equations  : 


HC1  +  NaOH  -  >-  NaCl  +  H2O  (or  HOH) 
HNO3  +  KOH  —  *•  KNO8  +  H2O 
H2SO4  +  Ca(OH)2  —  *-  CaS04  +  2  H2O 


In  all  such  reactions  the  acid  and  base  are  said  to  neu- 
tralize each  other,  and  the  process  is  termed  neutralization. 

Since  the  reactions  expressed  in  the  above  equations 
take  place  only  in  the  presence  of  water,  it  is  evident 
that  from  the  standpoint  of  the  ionization  theory  we  are 
really  dealing  here  with  ions  rather  than  with  compounds. 
Now,  experiments  show  that  hydrogen  ions  (H+)  and  hy- 
droxyl ions  (OH~)  cannot  exist  together  in  solution  to 
any  appreciable  extent,  but  at  once  combine  to  form  water  : 

H+  +  OH-—  >-H2O 

Considered  from  this  standpoint,  neutralization  consists  in 
the  union  of  the  hydrogen  ion  of  an  acid  with  the  hydroxyl 


ACIDS,  BASES,  AND  SALTS 


183 


ion  of  a  base  to  form  water.   To  express  this  fact  the  above 
reactions  should  be  written  as  follows: 


Na+,  C1- 


NO- 


Ca 


It  will  be  observed  that  in  neutralization  the  cation  of  the 
base  and  the  anion  of  the  acid  are  not  changed,  but  remain 
as  ions  in  the  solution.  If  now  the 
water  present  is  expelled  by  evap- 
oration, the  two  ions  unite  to  form 
a  compound  which,  with  few  ex- 
ceptions, is  a  solid  and  remains  as 
a  residue  in  the  dish.  Thus,  the 
ions  Na+  and  Cl~  unite  to  form 
the  white  solid  NaCl;  similarly, 
the  ions  K+  and  NOg~  unite  to  form 
potassium  nitrate  (KNOg),  while 
the  ions  Ca++  and  SO4 —  unite  to 
form  calcium  sulfate  (CaSO4). 

Neutralization  a  definite  act.  If 
two  solutions,  one  of  a  base  and 
the  other  of  an  acid,  are  prepared, 
experiment  has  shown  that  a  given 
volume  of  the  acid  will  invariably 
require  a  perfectly  definite  volume 
of  the  base  for  its  neutralization. 


I 


FIG.  83.   Two  burettes  used 
in  proving  that  neutraliza- 
tion is  a  definite  act 


The  experiment  is  most  easily  performed  with  the  aid  of 
burettes  (Fig.  83),  which  are  graduated  tubes  furnished  with 
a  stopcock  at  one  end.  The  one  is  filled  to  the  zero  mark  with 
the  acid  solution,  the  other  with  the  basic.  A  measured  vol- 
ume of  the  one  solution  is  drawn  off  into  a  small  beaker,  a 
few  drops  of  an  appropriate  indicator  added,  and  the  second 


184    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

solution  run  in  with  constant  stirring  until  the  indicator  just 
turns  color.  If  the  concentration  of  each  solution  is  accurately 
known,  it  is  easy  to  calculate,  from  the  volumes  required  for 
neutralization,  the  ratio  by  weight  between  the  acid  and  the 
base  taking  part  in  the  action.  Experiment  shows  that  this 
ratio  always  bears  a  simple  relation  to  that  between  the  molec- 
ular weights  of  the  reacting  substances.  Such  a  reaction  as  is 
indicated  in  the  equation 

Na+,  OH-  +  H+,  Cl- >-Na+,  Cl-  +  H20 

is  therefore  perfectly  definite,  and  if  we  know  the  weight  of 
an  acid  employed,  we  can  calculate  the  weight  of  the  base 
required  to  neutralize  the  acid  and  also  the  weight  of  the 
salt  formed. 

Heat  of  neutralization.  If  neutralization  is  due  to  the 
union  of  hydrogen  ions  with  hydroxyl  ions,  and  nothing 
more,  it  follows  that  when  a  given  weight  of  water  is 
formed  in  neutralization,  the  heat  set  free  should  always 
be  the  same,  no  matter  from  what  acid  and  base  the  two 
kinds  of  ions  have  been  supplied.  Careful  experiments 
have  shown  that  this  is  the  case,  provided  110  other  reac- 
tions take  place  at  the  same  time.  When  one  gram-molecule 
(18  g.)  of  water  is  formed  in  neutralization,  13,790  cal.  of 
heat  are  set  free.  This  is  represented  in  the  equations 

Na+,  OH-  -f  H+,  Cl-  — >•  Na+,  Cl~  +  H2O  +  13,790  cal. 
K+,  OH-  +  H+,  NO3-  — )-K+,  NO8-  +  H0O  -f  13,790  cal. 
Ca++,  (OH  )2+(H+)2,  SO4-- — >-Ca++;  SO4--+2H2O 
+  2  x  13,790  cal. 

Salts.    We  have  seen  in  the  process  of  neutralization 

(1)  that  the  cation  of  the  base  and  the  anion  of  the  acid 
are   not   changed,   but   remain   as   ions   in   the    solution ; 

(2)  that  if  the  water  present  is  expelled  by  evaporation, 


ACIDS,  BASES,  AND  SALTS  185 

the  two  ions  unite  to  form  a  compound.  The  compounds 
formed  in  this  way  by  the  union  of  the  cation  of  the  base 
and  the  anion  of  the  acid  are  termed  salts. 

The  following  equations  represent  the  reactions  taking 
place  in  the  formation  of  some  typical  salts. 

HC1  +  NaOH  — >•  NaCl  4-  H2O 
HC1  +  KOH  — *•  KC1  +  H20 

2HC1 4-  Ca(OH)2 *CaCl,  +  2  H0O 

HN08  +  KOH  — h  KN03  +  H26 
H2SO4  4-  Zn(OH)2  — >•  ZnSO4  +  2  H2O 

Relation  of  salts  to  acids.  The  compounds  represented 
by  the  formulas  NaCl,  KC1,  CaCl2,  KNO8,  ZnSO4  are  all 
salts.  If  we  examine  the  formulas  of  these  salts  carefully, 
we  shall  see  that  the  compounds  represented  by  these  for- 
mulas differ  from  the  acids  from  which  they  are  prepared 
in  that  an  atom  of  a  metal  has  been  substituted  for  one 
or  more  of  the  hydrogen  atoms  of  the  acid.  From  this 
standpoint,  a  salt  may  be  regarded  as  a  compound  which 
may  be  derived  from  an  acid  by  displacing  the  hydrogen  of 
the  acid  by  a  metal  It  is  customary  to  speak  of  the  salts 
derived  from  a  certain  acid  as  salts  of  that  acid.  Thus, 
NaCl,  KC1,  and  CaCl2  are  all  salts  of  hydrochloric  acid. 
Similarly,  KNO8  is  a  salt  of  nitric  acid,  while  ZnSO4  is  a 
salt  of  sulfuric  acid. 

Salts  are  also  formed  when  a  metal  dissolves  in  an  acid, 
as  illustrated  in  the  following  equations: 

Zn  4-  2  HC1 *•  ZnCl2  +  H2 

Fe  4-  H2S04  — *  FeSO4  +  Ha 

In  general,  we  might  expect  each  acid  to  form  as  many 
salts  as  there  are  metals,  and  in  most  cases  this  is  the  fact. 


186    AN  ELEMENTARY  STUDY  OF  CHEMISTEY 

lonization  of  salts.  All  salts  that  dissolve  in  water  ionize 
in  solution,  the  metal  of  the  salt  forming  the  cation,  while  the 
rest  of  the  molecule  forms  the  anion,  as  shown  in  the  follow- 
ing  equations:  NaC1_^Na+,  C1- 


ZnS04  -  ^Zn++,  SO4~ 
KNO3  —  >-K+,  NO8- 

It  will  be  seen  that  salts  differ  from  acids  and  bases  in 
that  the  solutions  have  no  common  ion.  It  follows  that 
salts  do  not  have  so  many  common  characteristics  as  do 
acids  and  bases. 

Normal  salts  ;  acid  salts.  It  is  evident  that  sodium 
hydroxide  can  act  upon  hydrochloric  acid  in  but  one 
proportion  : 


With  sulfuric  acid,  however,  the  reaction  may  take  place 
according  to  either  of  the  following  equations,  depending 
upon  the  relative  weights  of  sodium  hydroxide  used: 

NaOH  +  H2S04  —  +  NaHSO4  +  H2O 
2  NaOH  +  H2S04  —  *  Na2SO4  +  2  H2O 

It  will  be  observed  that  in  the  one  salt  (NaHSO4)  only 
a  portion  of  the  hydrogen  of  the  sulfuric  acid  has  been 
displaced  by  the  metal,  while  in  the  other  (Na2SO4)  all  of 
the  hydrogen  has  been  displaced.  To  distinguish  between 
these  two  classes  of  salts,  the  former  is  said  to  be  an  acid 
salt,  \Mhile  the  latter  is  termed  a  normal  salt. 

Acid  salts  when  dissolved  in  water  give  hydrogen  ions 
in  addition  to  the  ions  characteristic  of  salts  ;  thus, 

NaHSO4  —  >-Na+,  H+,  SO4 

They  have,  therefore,  the  properties  of  an  acid  as  well  as 
those  of  a  salt. 


ACIDS,  BASES,  AND  SALTS       187 

Basic  salts.  When  an  acid  such  as  hydrochloric  acid 
acts  upon  a  base  such  as  calcium  hydroxide,  the  reaction 
may  take  place  in  either  of  the  following  ways  : 

HC1  +  Ca(OH)2  —  »•  Ca(OH)Cl  +  H2O 
2  HC1  +  Ca(OH)2  —  >-CaCl2  +  2  H2O 


The  compound  Ca(OH)Cl  is  intermediate  between  a  salt 
and  a  base  ;  it  is  therefore  termed  a  basic  salt.  When  dis- 
solved in  water  it  gives  the  hydroxyl  ion  characteristic  of 
a  base  as  well  as  the  typical  salt  ions. 

Extent  of  ionization.  The  question  will  naturally  arise, 
When  an  acid,  base,  or  salt  dissolves  in  water,  do  all  the 
molecules  ionize  or  only  some  of  them  ?  The  experiments 
by  which  this  question  is  answered  cannot  be  described 
here.  It  has  been  found,  however,  that  only  a  fraction 
of  the  molecules  ionize.  The  percentage  which  ionizes  in 
a  given  case  depends  upon  several  conditions,  the  chief 
of  which  are  as  follows  : 

1.  The  concentration  of  the  solution.    In  concentrated  solu- 
tions only  a  very  small  percentage  of  the  molecules  ionize. 
As  the  solution  is  diluted  the  percentage  increases,  and  in 
dilute  solutions  it  may  be  very  large,  though  it  is  never 
complete  in  any  ordinary  solution.    Moreover,  an  acid  such 
as  sulfuric  acid  may  form  different  ions  according  to  the 
concentration  of  the  solution: 

H2SO4  (concentrated  solution)  -  >-  H+,  HSO4~ 

H2SO4  (dilute  solution)  -  >-H+,  H+,  SO," 

2.  The  nature  of  the  dissolved  compound.    At  equal  con- 
centrations, substances  differ  much  among  themselves  in 
percentage  of  ionization.     Most  salts   are   about  equally 
ionized.     Acids  and   bases,  on   the   contrary,   show  great 
differences,  some  being  freely  ionized,  others  very  slightly. 


188    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Strength  of  acids  and  bases.  Since  acid  and  basic  prop- 
erties are  due  to  hydrogen  ions  and  hydroxyl  ions,  the 
acid  or  base  which  will  produce  the  greatest  percentage 
of  these  ions  at  a  given  concentration  must  be  regarded  as 
the  strongest  representative  of  its  class.  The  acids  and 
bases  described  in  the  foregoing  paragraphs  are  all  quite 
strong.  In  10  per  cent  solutions  about  half  of  the  mole- 
cules are  dissociated  into  ions,  and  this  is  also  approxi- 
mately the  extent  to  which  most  salts  are  ionized  at  this 
same  concentration. 

Methods  of  expressing  reactions  between  compounds  in 
solution.  Chemical  equations  representing  reactions  between 
substances  in  solution  may  represent  the  details  of  the 
reaction,  or  they  may  simply  indicate  the  final  products 
formed.  Thus,  if  we  wish  to  call  attention  to  the  details 
of  the  reaction  between  sodium  hydroxide  and  hydrochloric 
acid  in  solution,  representing  the  ions  wliich  take  part  in 
the  reaction,  we  write  the  equation  as  follows: 

Na+,  OH-  -f  H+,  Cl-  -  >•  Na+,  Cl~  +  H2O 

If  we  wish  simply  to  represent  the  substances  taking  part 
in  the  reaction  and  the  final  products  formed,  we  write 
the  equation  thus  : 

NaOH  +  HC1  -  *  NaCl  +  H2O 

Similarly,  the  two  ways  of  expressing  the  reaction  be- 
tween zinc  and  hydrochloric  acid  are 

Zn  +  2   H+,  C1-   -  *-  Zn+  +,  (Cl-)g  +  H2 


Zn  +  2HCl 

Radicals.  We  have  seen  that  sulfuric  acid  as  well  as 
its  salts  all  contain  the  group  of  atoms  (SO4).  This  group 
acts  as  a  unit  and  forms  an  ion  when  the  acid  or  its  salts 


ACIDS,  BASES,  AND  SALTS  189 

are  dissolved  in  water.  Similarly,  nitric  acid  and  its  salts 
contain  the  group  (NOg),  while  all  bases  contain  the  group 
(OH).  All  such  groups  as  these,  which  act  as  a  unit  in 
chemical  reactions,  are  known  as  radicals.  We  may  define 
a  radical,  therefore,  as  a  group  of  elements  which  act  together 
as  a  unit  in  chemical  reactions.  Many  of  these  radicals 
have  been  given  special  names ;  for  example,  the  radical 
(OH)  is  known  as  the  hydroxyl  radical. 

Names  of  acids,  bases,  and  salts.  Since  acids,  bases,  and 
salts  are  so  intimately  related  to  each  other,  it  is  very 
advantageous  to  give  names  to  the  three  classes  in  accord- 
ance with  some  fixed  system.  The  system  universally 
adopted  is  as  follows : 

Naming  of  bases.  All  bases  are  called  hydroxides.  They 
are  distinguished  from  each  other  by  prefixing  the  name 
of  the  element  which  is  in  combination  with  the  hydroxyl 
group.  Examples:  sodium  hydroxide  (NaOH) ;  calcium 
hydroxide,  Ca(OH)0;  copper  hydroxide,  Cu(OH)2. 

Naming  of  acids.  The  method  of  naming  acids  depends 
upon  whether  the  acid  consists  of  two  elements  or  three. 

1.  Binary  acids.    Acids  containing  only  one  element  in 
addition  to  hydrogen  are  called  binary  acids.     They  are 
given  names  consisting  of  the  prefix  hydro-,  the  name  of  the 
second  element  present,  and  the  termination  -ic.   Examples : 
hydrochloric  acid  (HC1)  ;  hydrosulfuric  acid  (H0S). 

2.  Ternary  acids.   In  addition  to  the  two  elements  present 
in  binary  acids,  the  great  majority  of  acids  also  contain 
oxygen.    They  therefore  consist  of  three  elements  and  are 
called  ternary  acids.     It  usually  happens  that   the   same 
three  elements  can  unite  in  different  proportions  to  make 
several  different  acids.    The  most  familiar  one  of  these  is 
given  a  name  ending  in  the  suffix  -ic,  while  the  one  with 
less  oxygen  is  given  a  similar  name,  but  ending  in  the 


190    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

suffix  -ous.  Examples  :  nitric  acid  (HNO3)  ;  nitrous  acid 
(HNO2).  In  cases  where  more  than  two  acids  are  known, 
use  is  made  of  prefixes  in  addition  to  the  two  suffixes  '-ic 
and  -OILS.  Thus  the  prefix  per-  signifies  an  acid  still  richer 
in  oxygen  ;  the  prefix  hypo-  signifies  one  with  less  oxygen. 

Naming  of  salts.  A  salt  derived  from  a  binary  acid  is 
given  a  name  consisting  of  the  names  of  the  two  elements 
composing  it,  with  the  termination  -ide.  Example  :  sodium 
chloride  (NaCl).  All  other  binary  compounds  are  named 
hi  the  same  way. 

A  salt  of  a  ternary  acid  is  named  in  accordance  with 
the  acid  from  which  it  is  derived.  A  ternary  acid  whose 
name  has  the  termination  -ic  gives  a  salt  with  the  name 
ending  in  -ate,  while  an  acid  with  the  termination  -ous 
gives  a  salt  with  the  name  ending  in  -ite.  The  following 
table  will  make  the  application  of  these  principles  clear: 

ACID  FORMULA  SALT  FORMULA 

Hydrochloric      HC1  Sodium  chloride  NaCl 

Hypochlorous     HC1O  Sodium  hypochlorite     NaCIO 

Chlorous  HC1O2  Sodium  chlorite  NaClO2 

Chloric  HC1O3  Sodium  chlorate  NaClO3 

Perchloric  HC1O4  Sodium  perchlorate       NaClO4 

Electrochemical  series.  Upon  bringing  a  piece  of  zinc  into 
a  solution  of  an  acid,  zinc  passes  into  solution  and  hydrogen 
is  evolved  : 


In  like  manner,  when  zinc  is  placed  in  a  solution  of  a  salt 
of  copper,  such  as  the  sulfate  CuSO4,  zinc  passes  into  solu- 
tion, and  a  corresponding  quantity  of  copper  is  precipitated: 

Zn  +  CuSO4  -  >-  ZnSO4  +  Cu 

On  the  other  hand,  copper  has  no  effect  upon  a  solution 
of  zinc  sulfate. 


ACIDS,  BASES,  AND  SALTS 


'191 


It  has  been  found  to  be  possible  to  arrange  hydrogen 
and  the  metals  in  a  table  in  such  a  way  that  any  element 
in  the  list  will  displace  any  one  below  it  from  its  salts 
and  will  in  turn  be  displaced  from  its  salts  by  any  one 
above  it.  This  list  is  called  the  electrochemical  series  or 
the  displacement  series. 


ELECTROCHEMICAL  SERIES 


1.  Caesium 
2.  Rubidium 
3.  Potassium 
4.  Sodium 
5.  Lithium 

8.  Aluminium 
9.  Manganese 
10.  Zinc 
11.  Chromium 
12.  Cadmium 

15.  Nickel 
16.  Tin 
17.  Lead 
18.  Hydrogen 
19.  Arsenic 

22.  Bismuth 
23.  Mercury 
24.  Silver 
25.  Platinum 
26.  Gold 

6.  Calcium 


13.  Iron 


7.  Magnesium        14.  Cobalt 


20.  Copper 

21.  Antimony 


This  table  enables  us  to  foretell  many  reactions.  For 
example,  from  the  positions  of  the  two  metals  we  should 
expect  magnesium  to  displace  tin  from  its  salts : 

Mg  +  SnCl2 ^  MgCl2  +  Sn 

We  should  not,  however,  expect  iron  to  displace  aluminium. 
It  is  of  especial  interest  to  notice  the  position  of  hydro- 
gen in  the  series.  All  the  metals  above  it  will  evolve 
hydrogen  from  acids,  while  those  below  it  will  not.  In  the 
latter  case,  if  any  action  takes  place  it  must  be  preceded 
by  oxidation. 

From  the  standpoint  of  modern  electrical  theory  this  list 
really  represents  the  relative  ease  with  which  the  various 
atoms  give  up  one  or  more  electrons  to  form  ions.  Caesium, 
the  metal  going  into  solution  most  readily,  parts  with  an  elec- 
tron most  easily,  while  such  metals  as  gold  and  platinum 
retain  their  normal  number  of  electrons  most  tenaciously. 


192    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

EXERCISES 

1.  What  are  the   salts  of  nitric  acid  called?  the  salts  of  sul- 
furic  acid? 

2.  Name    the    acids   represented   by   the    following   formulas : 
H2SO4,  H2SO3,  H2SO2,  HNO8,  HNO2. 

3.  Give  the  name  of  each,  of  the  compounds  represented  by  the 
following   formulas  and  tell  to  which  group  of  compounds  each 
belongs  :  Mg(OH)2,  HBr,  NaBr,  H2SO3,  CaSO3,  PbSO4,  HI,  NaNO2, 
KNO3,  H2CO3,  Na2CO3,  FeCl8. 

4.  25  cc.  of  a  solution  containing  40  g.  of  sodium  hydroxide  per 
liter  was  found  to  neutralize  25  cc.  of  a  solution  of  hydrochloric 
acid.    What  was  the   strength  of  the  acid  solution  expressed  in 
grams  per  liter? 

5.  After  neutralizing  a  solution  of  sodium  hydroxide  with  nitric 
acid,  there  remained   after  evaporation  100  g.  of  sodium  nitrate. 
What  weight  of  each  compound  had  been  used? 

6.  A  solution  contains  18  g.  of  hydrogen  chloride  per  100  cc.    It 
required  25  cc.  of  this  solution  to  neutralize  30  cc.  of  a  solution  of 
sodium  hydroxide.    What  was  the  strength  of  the  sodium  hydroxide 
solution  expressed  in  grams  per  100  cc.  ? 

7.  When  perfectly  dry  sulfuric  acid  (hydrogen  sulfate)  is  treated 
with  perfectly   dry   sodium  hydroxide,  no  chemical  change  takes 
place.    Explain. 

8.  When  cold  concentrated  sulfuric  acid  is  added  to  zinc,  no 
change  takes  place.    Recall  the  action  of  dilute  sulfuric  acid  on  the 
same  metal.    How  do  you  account  for  the  difference  ? 

9.  A  solution  of  hydrogen  chloride  in  benzene  does  not  conduct 
the  electric  current.    When  this  solution  is  treated  with  zinc,  will 
hydrogen  be  evolved  ?    Explain. 

10.  (a)  Write  equation  for  preparation  of  hydrogen  from  zinc 
and  dilute  sulfuric  acid,    (ft)  Rewrite  the  same  equation  from  the 
standpoint  of  the  theory  of  electrolytic  dissociation,     (c)  Subtract 
the  common  ion  (SO4)  from  both  members  of  the  equation.    (V7)  From 
the  resulting  equation,  explain  in  what  the  preparation  of  hydrogen 
consists  when  examined  from  the  standpoint  of  this  theory. 

11.  In  the  same  manner  as  in  the  preceding  exercise,  explain  in 
what  the  action  of  sodium  on  water  to  give  hydrogen  consists. 


CHAPTER  XVII 
VALENCE 

Definition  of  valence.  A  comparison  of  the  composition 
of  the  compounds  of  hydrogen  with  the  other  elements 
brings  to  light  an  interesting  fact  illustrated  in  the  formulas 

HC1  H2O  H8N  H4C 

(hydrogen  chloride)  (water)  (ammonia)  (marsh  gas) 

It  will  be  seen  that  the  various  kinds  of  atoms  differ 
among  themselves  in  the  number  of  hydrogen  atoms  that 
they  are  able  to  hold  in  combination.  An  atom  of  chlorine 
combines  with  but  one  hydrogen  atom,  an  atom  of  oxygen 
with  two,  one  of  nitrogen  with  three,  and  one  of  carbon 
with  four.  It  is  convenient  to  have  a  name  to  designate 
that  property  of  an  element  that  determines  the  number  of 
hydrogen  atoms  that  its  atom  can  hold  in  combination.  It  is 
called  the  valence  of  an  element. 

Variety  in  valence.  One  atom  of  hydrogen  never  com- 
bines with  more  than  one  atom  of  any  other  kind,  so  that 
hydrogen  is  said  to  have  a  valence  of  1,  or  to  be  univalent. 
Other  elements,  such  as  chlorine,  iodine,  and  sodium,  which 
combine  with  hydrogen  atom  for  atom  (HC1,  HI,  HNa),  are 
likewise  said  to  be  univalent.  On  the  other  hand,  elements 
such  as  oxygen,  sulfur,  and  calcium,  one  atom  of  which 
combines  with  two  atoms  of  hydrogen  or  of  other  univalent 
elements,  are  said  to  be  bivalent.  Similarly,  we  have  triva- 
lent  elements,  such  as  nitrogen ;  and  quadrivalent  ones,  such 
as  carbon.  No  element  is  known  whose  valence  exceeds  8, 
and  with  most  elements  it  does  not  exceed  4, 


194    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

How  elements  of  different  valences  combine.  If  the  valence 
of  two  elements  is  the  same,  we  may  expect  them  to  com- 
bine atom  for  atom.  Thus  the  univalent  elements,  hydro- 
gen and  chlorine,  combine  to  form  the  compound  HC1 ; 
the  bivalent  elements,  calcium  and  oxygen,  to  form  the 
compound  CaO ;  the  trivalent  elements,  nitrogen  and  alu- 
minium, to  form  A1N.  When  the  valences  of  two  elements 
are  different,  the  ratio  between  the  atoms  is  not  so  simple. 
The  sum  of  the  valences  of  the  one  kind  of  atoms  in  the 
molecule  must  equal  the  sum  of  the  valences  of  the  other 
kind  of  atoms  ;  and  this  will  require  the  two  kinds  of  atoms 
to  be  present  in  unequal  numbers.  For  example,  bivalent 
calcium  and  univalent  chlorine  will  give  the  compound 
CaCl2 ;  trivalent  aluminium  and  univalent  chlorine  will  give 
A1C18 ;  trivalent  iron  and  bivalent  oxygen  will  give  Fe2Og ; 
quadrivalent  carbon  and  bivalent  oxygen  will  give  CO2. 

Hydrogen  and  oxygen  the  standards  for  valence.  It  will 
be  noticed  that  the  fact  that  two  elements  combine  in  the 
ratio  of  1  atom  each,  as  in  CaO,  does  not  necessarily  indi- 
cate that  these  elements  are  univalent.  The  hydrogen  atom 
is  taken  as  the  standard,  and  with  it  calcium  forms  the 
compound  CaH2  and  is  therefore  bivalent.  Consequently  in 
the  compound  CaO  calcium  as  well  as  oxygen  is  bivalent. 

Since  the  oxides  of  the  elements  are  much  better  known 
than  the  hydrides,  we  often  determine  the  valence  of  an 
element  by  examining  the  formula  of  its  oxide  rather  than 
that  of  its  hydride.  As  is  indicated  by  the  formula  H2O, 
oxygen  is  bivalent.  Consequently  from  the  formulas  Na2O, 
CaO,  A12O3  we  infer  that  sodium  (Na)  is  univalent ;  cal- 
cium (Ca)  is  bivalent ;  and  aluminium  (Al)  is  trivalent. 

Probable  formulas  from  known  valences.  If  we  know 
the  valences  of  two  elements,  we  can  determine  the  probable 
formula  which  at  least  one  of  their  compounds  will  have. 


VALENCE  195 

Thus,  if  we  know  that  bromine  (Br)  is  univalent  (as  is  shown 
by  the  formula  HBr)  and  that  aluminium  is  trivalent,  we 
may  expect  that  the  formula  of  aluminium  bromide  will 
be  AlBrs;  if  phosphorus  (P)  is  pentavalent  and  oxygen 
is  bivalent,  phosphoric  oxide  will  have  the  formula  P2O5. 

Valence  of  radicals.  We  have  found  that  sometimes  a 
group  of  atoms  or  a  radical  takes  part  in  a  reaction  very 
much  as  though  it  were  an  atom  (p.  188).  It  is  very 
convenient  to  think  of  such  a  radical  as  having  a  valence 
like  an  element.  Accordingly,  from  the  formula  H(OH) 
or  Na(OH)  we  argue  that  the  hydroxyl  radical  is  uni- 
valent. Since  calcium  is  bivalent,  the  formula  of  calcium 
hydroxide  will  therefore  be  Ca(OH)2.  From  the  formula 
H2(SO4)  we  argue  that  the  radical  (SO4)  is  bivalent ; 
calcium  sulfate  should  therefore  have  the  formula  CaSO4 
and  aluminium  sulfate  the  formula  A12(SO4)3. 

The  replacing  power  of  atoms  and  radicals.  Just  as  ele- 
ments having  the  same  valence  combine  with  each  other 
atom  for  atom,  so,  if  they  replace  each  other  in  a  chemical 
reaction,  they  will  do  so  in  the  same  ratio.  Thus,  one 
atom  of  bivalent  zinc  displaces  two  atoms  of  univalent 
hydrogen,  as  is  shown  in  the  following  equations: 

Zn  +  H0SO4 *-  ZnSO4  +  H2 

Zn  (OH)2  +  H2SO4 >•  ZnSO4  +  2  H2O 

Similarly,  one  atom  of  bivalent  calcium  displaces  one  atom 
of  bivalent  zinc : 

CaCl2  +  ZnSO4  — *•  CaSO4  +  ZnCl2 

Since  many  reactions,  like  those  above,  consist  in  the 
interchange  of  two  elements,  it  is  evident  that  a  knowledge 
of  the  valence  of  the  elements  will  assist  us  in  writing  the 
equations  for  the  reactions. 


196    AX  ELEMENTARY  STUDY  OF  CHEMISTRY 

Variable  valence.  It  often  happens  that  two  given  ele- 
ments form  more  than  one  compound,  and  in  such  cases 
at  least  one  of  the  elements  must  have  more  than  one 
valence.  We  consider  oxygen  to  be  almost  always  bivalent, 
so  that  in  the  oxide  CO  carbon  must  be  bivalent,  while  in 
the  oxide  CO2  it  is  quadrivalent.  Similarly,  in  the  oxides 
FeO  and  Fe2O3  iron  is  bivalent  and  trivalent  respectively. 
In  general,  it  is  true  that  each  element  has  one  valence 
that  is  much  more  frequently  exerted  than  any  other,  so 
that  we  can  think  of  a  given  element  as  having  in  the  main 
a  certain  valence  and  less  frequently  some  other  valence. 

Valence  and  structure  of  molecules.  When  a  compound 
contains  more  than  two  elements,  as  H2SO4,  it  is  not 
practicable  to  determine  the  valence  of  the  several  atoms 
from  the  formula,  for  we  cannot  tell  in  what  way  the  ele- 
ments are  combined.  If  we  make  use  of  lines  drawn  between 
the  symbols  of  the  atoms  to  indicate  valence  as  well  as  the 
relation  of  the  atoms,  the  formula  of  water  might  be  written 
H—  O—  H.  This  would  mean  that  each  hydrogen  atom  is 
univalent,  while  the  oxygen  is  bivalent  ;  and  also  that  each 
hydrogen  atom  is  in  combination  with  the  oxygen  atom. 
Similarly,  we  might  write  the  formula  for  sulfuric  acid  in 
a  number  of  different  ways,  all  of  which  would  show  the 
hydrogen  to  be  univalent  and  the  oxygen  bivalent,  yet 
the  valence  of  sulfur  might  be  different  in  each  formula. 
In  the  following  formulas  the  possible  valences  of  sulfur 
are  indicated  by  the  numerals  placed  below  the  symbols: 


n 

H-0-=0          H-O-o-O          H-O-O- 


We  should  have  to  know  which  of  these  formulas  expresses 
the  truth  before  we  could  infer  the  valence  of  sulfur  in 


VALENCE  197 

the  compound.  Experiments  indicate  that  the  arrangement 
of  the  atoms  represented  in  formula  (jB)  is  the  correct 
one,  so  that  the  sulfur  in  sulfuric  acid  is  hexavalent. 
Formulas  like  the  ones  just  given,  which  represent  the 
arrangement  of  the  atoms  in  the  molecule,  are  called 
structural  formulas. 

Even  in  compounds  containing  but  two  elements  it  is  not 
always  possible  to  draw  correct  conclusions  from  the  mere  for- 
mula. Thus,  in  the  formula  for  magnetic  oxide  of  iron  (Fe804), 
if  we  assume  oxygen  to  have  its  usual  valence  of  2,  the  four 
atoms  will  give  a  total  of  8.  If  the  three  atoms  of  iron  are 
all  directly  united  with  oxygen  in  the  same  way,  then  each 
atom  of  iron  would  have  a  valence  of  2§.  It  can  be  shown, 
however,  that  the  structural  formula  of  the  oxide  is  as  follows  : 

0  =  Fe-0.  F 
0  =  Fe-0> 

It  will  be  seen  that  two  of  the  atoms  of  iron  in  the  oxide 
are  trivalent,  while  the  other  one  is  bivalent. 

The  cause  of  valence.  If  it  be  asked,  Why  do  atoms 
differ  in  valence  so  that  one  atom  of  chlorine  can  combine 
with  but  one  of  hydrogen,  while  one  of  oxygen  can  com- 
bine with  two?  the  answer  will  have  to  be  that  we  do 
not  know.  It  seems  probable,  however,  that  the  cause  is  to 
be  found  in  the  capacity  of  atoms  of  the  various  elements 
to  take  up  different  charges  of  electricity. 

We  have  seen  that  the  electrical  charges  on  the  ions 
of  an  electrolyte  in  solution  are  always  multiples  of  a 
single  unit  charge.  Thus,  with  hydrogen  chloride  the  ions 
formed  have  each  a  unit  charge,  one  ion  being  plus  (+)» 
the  other  minus  (— )  : 

HC1 — »- 


198    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

This  corresponds  to  the  fact  that  hydrogen  and  chlorine 
are  univalent.    With  calcium  sulfate  we  have  the  equation 


CaSO 


Ca 


The  double  charge  corresponds  to  the  bivalency  of  calcium 
and  of  the  radical  (SO4). 

It  seems  to  be  probable  that  a  similar  explanation  holds 
good  for  valency  in  general.  If  a  given  atom  tends  to 
take  up  three  positive  charges,  it  will  have  a  valence  of 
three,  like  aluminium,  and  will  tend  to  combine  with  atoms 
or  radicals  having  negative  charges.  If  it  takes  up  two 
negative  charges,  it  will  have  a  valence  of  two,  like  oxygen, 
and  will  tend  to  combine  with  atoms  or  radicals  having 
positive  charges.  These  facts  lead  us  to  speak  of  elements 
as  either  electropositive  or  electronegative.  When  charged 
atoms  unite  they  form  molecules  consisting  of  the  two 
kinds  of  atoms,  and  in  such  ratios  that  the  positive  charge 
of  the  one  just  neutralizes  the  negative  charge  of  the  other. 

Table  of  valences.  It  will  be  convenient  for  reference  to 
tabulate  the  valences  of  the  most  familiar  elements  and 
radicals  and  the  number  of  charges  they  carry  as  ions. 
Some  elements  will  be  found  in  several  different  lines, 
since  they  have  more  than  one  valence. 


TABLE  OF  VALENCES 


VALENCE 

POSITIVE  IONS 

NEGATIVE  IONS 

NOT  IONS 

1 

H,  Na,  K,  Ag,  NH4 

Cl,  Br,  I,  OH,  N03 

2 

Ca,  Ba,  Mg,  Zn,  Hg, 

Cu,  Fe,  Sn 

S,  S04,  C03 

0 

3 

Al,  Bi,  Sb,  Fe 

P04 

N,  P 

4 

Sn 

SiO4 

C,  Si,  S 

5 

N,  P,  As,  Sb 

6 

S 

VALENCE  199 

EXERCISES 

1.  In    the   compounds   whose    formulas    follow,  the   valence  of 
one  of  the  elements  is  indicated  by  a  figure  placed  over  the  sym- 
bol.   What  is  the  valence  of  each  of  the  other  elements  or  radicals 
present  in  the  compounds? 

222'  2  1  1 

MgBr2;  CaO ;  Ca(OH)2;  Ba8(PO4)2 ;  KNO8;  H2O 

Verify  your  results  by  reference  to  the  table. 

2.  Complete  and  balance  the  following  equations,  assuming  that 
in  each  instance  the  sodium  or  the  calcium  in  the  first  compound 
changes  place  with  the  hydrogen  of  the  second  : 


NaOH  +  HC1 >-  —    —  + 

NaOH  +  H2S04 >-—     -  + 

NaOH  +  H3PO4 >-  —     -  + 

Ca(OH)2  +  HC1 ) + 

Ca(OH)2  +  H2S04 ) -  + 

Ca(OH)0  +  H3P04 ) + 


3.  Sodium,  calcium,  and  aluminium  have  valences  of  1,  2,  and  3 
respectively ;    write  the  formulas  of  their  chlorides,  sulf ates,  and 
phosphates  (phosphoric  acid  =  H3PO4),  on  the  supposition  that  they 
form  salts  having  the  normal  composition. 

4.  Iron  forms  one  series  of  salts  in  which  it  has  a  valence  of  2, 
and  another  series  in  which  it  has  a  valence  of  3  ;  write  the  formulas 
for  the  two   chlorides  of  iron,  also  for  the  two   sulfates,  on  the 
supposition  that  these  have  the  normal  composition. 

5.  Silver  acts  as  a  univalent  element  and  calcium  as  a  bivalent 
element  in  the  formation  of  their  respective  nitrates  and  chlorides. 

(a)  Write   the    formula   for  silver   nitrate;    for  calcium   chloride. 

(b)  When  solutions  of  these  two  salts  are  mixed,  the  two  metals, 
silver  and  calcium,  exchange  places;    write  the   equation  for  the 
reaction. 

6.  Antimony  acts  as  a  trivalent  element  in  the  formation  of  a 
chloride,   (a)  What  is  the  formula  for  antimony  chloride  ?    (ft)  When 
hydrogen  sulfide  (H2S)  is  passed  into  a  solution  of  this  chloride,  the 
hydrogen  and  antimony  exchange  places ;  write  the  equation  for 
the  reaction. 


CHAPTER  XVIII 
COMPOUNDS  OF  NITROGEN 

Occurrence.  Large  quantities  of  nitrogen  occur  in  the 
atmosphere,  and  it  is  substantially  all  in  the  free  state.  In 
the  materials  composing  the  earth's  crust,  on  the  other  hand, 
there  occur  in  certain  localities  considerable  deposits  of  com- 
pounds of  nitrogen,  especially  of  sodium  nitrate  (NaNO3). 
Moreover,  such  compounds  are  present,  at  least  in  small 
quantities,  in  all  productive  soils.  From  these  soils  the  nitro- 
gen is  taken  up  by  plants  and  built  into  complex  compounds. 
Animals  feeding  on  these  plants  assimilate  the  nitrogenous 
matter,  which  becomes  an  essential  part  of  the  animal  tissue. 
In  both  plants  and  animals  the  nitrogen  is  present  chiefly 
in  the  form  of  proteid  matter,  which  consists  of  complex 
compounds  containing  the  elements  carbon,  nitrogen,  oxy- 
gen, and  hydrogen,  and  sometimes  phosphorus  and  sulfur. 

The  unstable  character  of  compounds  of  nitrogen.  Experi- 
ment shows  that  the  molecule  of  nitrogen  has  the  formula 
N2  and  that  the  element  is  very  inactive  at  ordinary  tem- 
peratures. This  inactivity  seems  to  be  partly  due  to  the 
fact  that  the  nitrogen  molecule  is  very  stable  and  that  a 
good  deal  of  energy  is  required  to  separate  it  into  its 
atoms,  which  must  be  done  before  it  can  enter  into  com- 
bination with  other  elements.  On  the  other  hand,  when 
nitrogen  occurs  as  a  constituent  of  a  compound,  the  nitro- 
gen atoms  tend  to  leave  the  compound,  with  liberation  of 
energy,  and  form  stable  nitrogen  molecules.  As  a  result 
of  this  tendency  compounds  containing  nitrogen  are  apt  to 


COMPOUNDS  OF  NITROGEN  201 

be  unstable.  It  is  largely  due  to  the  unstable  character  of 
certain  nitrogenous  compounds  that  they  are  so  extensively 
used  as  a  constituent  of  explosives. 

While  a  great  many  compounds  of  nitrogen  are  known, 
it  is  desirable  at  this  time  to  become  acquainted  with  only 
some  of  the  simple  ones,  especially  those  which  nitrogen 
forms  with  hydrogen  and  with  oxygen. 

COMPOUNDS  OF  NITEOGEN  WITH  HYDROGEN 

Nitrogen  forms  three  simple  compounds  with  hydrogen, 
the  names  and  formulas  of  which  are  as  follows :  ammonia 
(NH3),  hydrazine  (N2H4),  and  hydronitric  acid  (HNg). 
Of  these  ammonia  is  by  far  the  most  important. 

Ammonia.  Inasmuch  as  ammonia  is  formed  in  certain 
natural  processes  which  are  constantly  taking  place  about 
us,  such  as  the  decay  of  nitrogenous  organic  matter,  it  is 
easy  to  understand  why  this  compound  has  been  known 
for  so  long  a  time.  It  was  originally  prepared  by  heating 
such  tissues  as  the  hoofs  and  horns  of  animals,  and  the 
aqueous  solution  of  the  gas  so  obtained  was  termed  spirits 
of  hartshorn.  The  pure  gas  itself  was  first  prepared  by 
Priestley,  in  1774,  and  its  composition  was  determined 
soon  after  by  the  French  chemist  Berthollet. 

Preparation  of  ammonia.  The  principal  methods  for 
preparing  ammonia  are  as  follows: 

1.  Laboratory  method.  In  the  laboratory,  ammonia  is 
usually  prepared  from  ammonium  chloride  (NH4C1),  a 
white  solid  obtained  in  the  manufacture  of  coal  gas.  When 
a  mixture  of  ammonium  chloride  and  sodium  hydroxide  is 
heated,  the  ammonium  radical  (NH4)  and  sodium  change 
places,  as  represented  in  the  following  equation: 

NH4C1  +  NaOH >-  NaCl  +  NH4OH 


202    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


The  resulting  ammonium  hydroxide  (NH4OH)  is  unstable 
and,  as  fast  as  it  is  formed,  breaks  down  into  water  and 
ammonia : 


NH4OH 


Calcium  hydroxide,  Ca(OH)2,  is  frequently  used  in  place 
of  the  more  expensive  sodium  hydroxide : 


2NHCl 


Ca(OH), 
2NHOH 


CaCl2+2NH4OH 
2H20  +  2NH8 


The  ammonium  chloride  and  calcium  hydroxide  are  mixed 
together  and  placed  in  a  flask  A,  arranged  as  shown  in  Fig.  84. 

The  mixture  is  gently 
warmed,  when  ammonia 
is  evolved  as  a  gas  andj 
being  much  lighter  than 
air,  is  collected  in  B  by 
displacement  of  air,  as 
shown  in  the  diagram. 

2.  Commercial  prep- 
aration. Ammonia  is 
obtained  commercially 
in  the  process  of  man- 
ufacturing coal  gas 
and  coke.  Certain 
grades  of  soft  coal 
are  best  adapted  for 
this  purpose.  Such 
coal  contains,  in  addition  to  carbon,  about  1  per  cent  of 
nitrogen  and  7  per  cent  of  hydrogen,  as  well  as  small  per- 
centages of  other  elements.  When  such  coal  is  heated  in 
retorts  from  which  the  air  is  excluded  (p.  306),  complicated 
changes  take  place,  resulting  not  only  in  the  formation  of 
the  combustible  gases  which  constitute  coal  gas  but  also 


FIG.  84.    The  preparation  and  collection  of 
ammonia  in  the  laboratory 


COMPOUNDS  OF  NITROGEN  203 

of  ammonia  and  many  other  valuable  products.  From  25  to 
50  per  cent  of  the  nitrogen  present  in  the  coal  is  converted 
into  ammonia.  The  volatile  matter  expelled  from  the  coal 
is  passed  through  water,  which  absorbs  the  ammonia,  to- 
gether with  certain  other  compounds,  forming  a  solution 
known  as  the  ammoniacal  liquor.  When  this  liquor  is  heated 
with  slaked  lime,  ammonia  is  evolved  and  is  absorbed  either 
in  water  or  in  a  dilute  solution  of  an  acid.  • 

3.  The  synthetic  method  (the  Haber  process).  Ammonia  may 
be  formed  by  the  direct  union  of  nitrogen  and  hydrogen. 
The  best  yield  is  obtained  when  a  mixture  of  the  gases, 
subjected  to  a  pressure  of  200  atmospheres,  is  heated  to 
about  500°  in  contact  with  finely  divided  iron,  which  acts 
as  a  catalytic  agent.  Under  these  conditions,  a  small  per- 
centage of  ammonia  is  formed.  The  gases  are  then  passed 
through  water,  which  dissolves  the  ammonia,  while  the 
nitrogen  and  hydrogen  mixture  is  again  conducted  over  the 
heated  iron.  The  process  thus  becomes  continuous,  more 
nitrogen  and  hydrogen  being  introduced  as  needed.  This 
process,  known  as  the  Haber  process  from  the  German 
chemist  who  devised  it,  is  now  used  to  a  certain  extent  in 
Germany  for  the  production  of  ammonia  on  a  large  scale. 

Ammonia  is  likewise  formed  when  an  electric  discharge 
is  passed  through  a  mixture  of  nitrogen  and  hydrogen, 
but  the  yield  is  very  small. 

Properties.  Ammonia  is  a  colorless  gas  having  a  strong, 
suffocating  odor.  Under  standard  conditions  1  1.  of  the 
pure  gas  weighs  0.7708  g.,  being  0.59  times  as  heavy  as 
air.  The  gas  is  easily  condensed  to  a  colorless  liquid  boil- 
ing at  —  33.5°,  and  in  this  form  is  an  article  of  commerce. 
Liquid  ammonia,  like  water,  is  not  only  an  excellent  solvent 
but  also  a  highly  ionizing  one.  Ammonia  can  be  obtained 
in  the  form  of  a  snowlike  solid  melting  at  —  75.5°. 


204    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Ammonia  is  extremely  soluble  in  water,  1  1.  of  water 
at  0°  and  760  mm.  pressure  dissolving  1298.9  1.  of  the  gas, 
and  710  1.  at  20°.  In  dissolving  such  large  volumes  of  the 
gas  the  water  expands  considerably,  so  that  the  density  of 
the  solution  is  less  than  that  of  water.  The  most  concen- 
trated commercial  solutions  have  a  density  of  0.88,  and 
contain  35.6  per  cent  by  weight  of  the  gas. 

Chemical  conduct.  At  ordinary  temperatures  ammonia 
is  a  stable  compound.  When  heated  to  high  temperatures, 
however,  or  when  subjected  to  the  action  of  an  electric 
discharge,  it  is  decomposed  into  its  elements.  It  will  burn 
in  an  atmosphere  of  oxygen,  but  not  in  air  unless  heat  is 
continuously  applied. 

The  hydrogen  of  ammonia  can  be  displaced  by  metals 
forming  nitrides.  Thus,  magnesium  forms  the  white  solid 
magnesium  nitride  (MggN2)  : 

3  Mg  +  2NH3  —  >-Mg3N2  +  3  Ha 

It  will  be  observed  that  the  decomposition  of  ammonia  by  heat- 
ing is  just  the  reverse  of  the  formation  of  ammonia  in  the  Haber 
process.  These  reactions  will  be  discussed  in  Chapter  XIX. 

Action  of  ammonia  upon  water  ;  ammonium  hydroxide. 
The  solution  of  ammonia  in  water,  called  aqua  ammonia, 
is  found  to  have  strong  basic  properties.  It  turns  red 
litmus  blue  ;  it  feels  soapy  to  the  touch  ;  it  neutralizes 
acids,  forming  salts  with  them.  It  seems  certain,  there- 
fore, that  when  ammonia  dissolves  in  water,  a  portion  of 
the  gas  combines  chemically  with  water  according  to  the 
equation 


It  is  the  substance  NH4OH,  called  ammonium  hydroxide, 
which  has  the  basic  properties,  dissociating  into  the  ions 
NH4+  and  OH~.  At  ordinary  temperatures  the  separation 


COMPOUNDS  OF  NITKOGEN  205 

of  the  pure  hydroxide  from  its  solutions  has  not  been 
accomplished,  for  as  the  solution  becomes  concentrated 
the  compound  decomposes  again  into  ammonia  and  water. 
Ammonium  salts.  When  an  acid  is  added  to  aqua  am- 
monia, the  ammonium  hydroxide  and  the  acid  neutralize 
each  other  according  to  the  following  equations: 

NH4OH  +  HC1  — >•  NH4C1  +  HaO 
2  NH4OH  +  H2SO4  — +  (NH4)2SO4  +  2  H2O 

Upon  evaporation  the  resulting  salts,  NH4C1  and  (NH4)2SO4, 
are  obtained  in  the  form  of  white  solids.  In  these  salts 
the  radical  NH4  plays  the  part  of  a  metal.  On  this  account 
the  name  ammonium  is  given  to  it,  since  the  names  of 
most  of  the  metals  end  in  ium.  Salts  containing  this 
radical  are  known  as  ammonium  salts.  Some  of  the  most 
common  of  these  salts  are  ammonium  chloride  (NH4C1), 
ammonium  sulfate,  (NH4)0SO4,  and  ammonium  nitrate 
(NH4N03). 

Uses  of  ammonia.  Large  quantities  of  ammonia  are  used 
in  the  manufacture  of  aqua  ammonia,  as  well  as  in  the 
formation  of  ammonium  compounds,  such  as  ammonium 
chloride  and  ammonium  sulfate.  In  the  liquid  state  it  is 
also  used  extensively  in  the  manufacture  of  artificial  ice 
(p.  110).  Its  use  for  this  purpose  is  based  on  the  facts 
that  the  gas  is  easily  liquefied  by  pressure  and  that  the 
resulting  liquid  has  a  relatively  high  heat  of  vaporization. 

Composition  of  ammonia.  That  ammonia  is  a  compound 
of  nitrogen  and  hydrogen  is  proved  by  the  fact  that  it  may 
be  formed  by  the  direct  union  of  the  two  elements  (see 
method  of  preparation).  The  quantitative  composition  of 
the  compound  may  be  determined  by  taking  advantage 
of  certain  reactions  which  make  it  possible  to  liberate  the 
nitrogen  as  well  as  the  hydrogen  from  any  definite  volume 


206    AN  ELEMENTAKY  STUDY  OF  CHEMISTRY 

of  ammonia.  By  measuring  the  volumes  of  the  gases  so 
liberated  one  can  compare  them  not  only  with  each  other 
but  also  with  the  volume  of  the  ammonia  from  which  they 
were  derived.  In  this  way  it  has  been  proved  that  2  vol- 
umes of  ammonia  yield  on  decomposition  1  volume  of 
nitrogen  and  3  volumes  of  hydrogen,  as  expressed  graphb 
cally  in  the  following  equation: 


N2  H2  +  H2  +  H2 

Hydrazine  (N2H4).  This  is  a  colorless  liquid  boiling  at  113.5° 
Like  ammonia,  it  combines  with  water  to  form  a  base,  from 
which  salts  can  be  prepared  by  the  action  of  acids. 

Hydronitric  acid  (HN3).  This  acid  is  a  colorless  liquid  of  dis- 
agreeable odor.  It  boils  at  37°  and  is  violently  explosive, 
decomposing  into  its  constituent  elements  with  the  liberation 
of  considerable  heat.  The  salts  of  hydronitric  acid  are  solids. 
Some  of  them  are  violently  explosive  and  are  used  in  the 
manufacture  of  explosives. 

COMPOUNDS  OF  NITROGEN  WITH  HYDROGEN  AND  OXYGEN 

Nitrogen  forms  a  number  of  compounds  with  hydrogen 
and  oxygen,  the  most  important  of  which  are  the  two 
acids,  nitric  acid  (HNO3)  and  nitrous  acid  (HNO2). 

Nitric  acid.  Nitric  acid  was  well  known  to  the  alche- 
mists, being  first  prepared  by  the  Egyptians.  In  the  ninth 
century  the  alchemist  Geber  prepared  it  from  saltpeter 
(KNOg)  by  a  process  somewhat  similar  to  that  used  at 
the  present  time,  and  the  Germans  still  call  it  salpeter- 
saure.  The  composition  of  the  acid  was  first  determined 
by  Lavoisier  and  Priestley. 

Because  of  its  great  activity  nitric  acid  does  not  occur 
free  in  nature,  but  a  number  of  its  salts  are  found  in 


COMPOUNDS  OF  NITROGEN 


207 


considerable  quantities.  The  most  abundant  of  these  is 
sodium  nitrate  (NaNO8),  which  is  found  in  large  quan- 
tities in  Chile  and  is  known  as  Chile  saltpeter. 

Preparation  of  nitric  acid.  Nitric  acid  can  be  prepared 
by  a  number  of  different  methods,  the  most  important  of 
which  are  the  following: 

1 .  Preparation  from  sodium  nitrate.  When  sodium  nitrate 
is  treated  with  concentrated  cold  sulfuric  acid,  no  chemical 
action  seems  to  take  place.  If,  however,  the  mixture  is 
placed  in  a  retort 
A  (Fig.  85)  and  a 
gentle  heat  applied, 
nitric  acid  is  given 
off  as  a  vapor  and 
may  be  condensed 
to  a  liquid  by  con- 
ducting the  vapor 
into  a  tube  B  sur- 
rounded by  ice  water, 
as  shown  in  Fig.  85. 
An  examination  of 
the  liquid  left  in  the 


FIG.  85.    The  preparation  of  nitric  acid  in 
the  laboratory 


retort  shows  that  it  contains  acid  sodium  sulfate  (NaHSO4), 
half  of  the  hydrogen  of  sulfuric  acid  having  been  replaced 
by  sodium.  The  equation  for  the  reaction  is  as  follows : 


0. — >-NaHSC>  +  HNOa 

•  »          4  »  3 

If  a  smaller  quantity  of  sulfuric  acid  is  taken  and  the  mix- 
ture is  heated  to  a  high  temperature.,  normal  sodium  sulfate  is 
formed : 

2  NaN08  +  H2S04 >-  Na3SO4  +  2  HN08 

In  this  case,  however,  the  higher  temperature  required  decom- 
poses a  part  of  the  nitric  acid,  so  the  process  is  not  economical. 


208    AN  ELEMENTARY  STUDY  OF  CHEMISTKY 


The  commercial  preparation  of  nitric  acid.  Fig.  86  illustrates 
a  form  of  apparatus  used  in  the  preparation  of  nitric  acid  on 
a  large  scale.  Sodium  nitrate  and  sulfuric  acid  are  heated  in 
the  iron  retort  A.  The  resulting  acid  vapors  pass  in  the  direc- 
tion indicated  by  the  arrows,  and  are  condensed  in  the  glass 
tubes  B,  which  are  covered  with  cloth  kept  cool  by  streams 
of  water.  These  tubes  are  inclined  so  that  the  liquid  resulting 
from  the  condensation  of  the  vapors  runs  back  into  C  and 

is  drawn  off  into 
large  vessels  (D). 


2.  Preparation 
from  air.  When 
an  electric  dis- 
charge takes  place 
through  a  mix- 
ture of  oxygen 
and  nitrogen  (air), 
a  small  percen- 
tage of  oxides 
of  nitrogen  is 
formed.  This  can 
be  increased  by 

having  the  mixture  pass  through  an  electric  arc  which  has 
been  drawn  out  to  a  great  size  by  magnets  (Fig.  87).  The 
oxides  so  obtained  combine  with  water  to  form  dilute  nitric 
acid.  This  method  for  preparing  nitric  acid  (known  as 
the  Birkeland  and  Eyde  process)  has  come  into  extensive 
use  in  recent  years  in  Norway,  since  the  necessary  elec- 
trical energy  can  be  generated  at  a  very  low  cost  by  the 
waterfalls  abounding  in  that  country.  The  dilute  nitric 
acid  obtained  is  neutralized  with  lime  (CaO),  ,and  the 
resulting  calcium  nitrate  is  sold  for  use  as  a  fertilizer 
under  the  name  air  saltpeter. 


•    FIG. 


A  commercial  still  for  the  production  of 
concentrated  nitric  acid 


COMPOUNDS  OF  NITROGEN  209 

3.  Preparation  from  ammonia.  Nitric  acid  is  also  formed 
by  the  oxidation  of  ammonia.  The  process  consists  in 
heating  a  mixture  of  ammonia  and  air  in  contact  with 
platinum  or  some  other  catalyzer. 

The  preparation  of  nitric  acid  from  sodium  nitrate  is,  under 
normal  conditions,  the  most  economical  process.  The  methods 
used  in  its  preparation  from  air  and  from  ammonia,  however, 
are  coming  into  use  and  undoubtedly  will  displace  the  sodium 
nitrate  method  as  the  supply  of  this  salt  becomes  exhausted. 
During  the  great  war  the  Central  Powers  could  not  obtain 
sodium  nitrate  and  hence  prepared  the  enormous  supplies  of 
nitric  acid  required  in  the  manu- 
facture of  explosives  chiefly  from 
ammonia. 


Properties  of  nitric  acid.    Pure 
nitric   acid  (hydrogen  nitrate)  is 
a  colorless  liquid  which  boils  at 
about  86°  and  has  a  density  of     FIG.  87.  Form  of  the  electric 
.,  r/>      rri  ,    j         -j       P     arc  employed  in  the  Birke- 

1.56.    The   concentrated   acid   of        iand  and  Eyde  process 
commerce  contains  about  68  per 

cent  of  the  acid,  the  remainder  being  water.  Such  a  mix- 
ture has  a  density  of  1.4.  The  concentrated  acid  fumes 
somewhat  in  moist  air,  and  has  a  sharp,  choking  odor. 

Chemical  conduct.  The  most  important  chemical  reac- 
tions of  nitric  acid  are  the  following: 

1.  Acid  action.  Nitric  acid  has  all  the  characteristics  of 
a  strong  acid.  It  changes  blue  litmus  red,  and  has  a  sour 
taste  in  dilute  solutions.  It  gives  the  ions  H+  and  NO8~ 
in  solution,  and  neutralizes  bases,  forming  salts.  It  also 
acts  upon  the  oxides  of  most  metals,  forming  a  salt  and 
water;  thus, 

CuO  +  2  HN03 — >^Cu(NO8)2  +  H2O 


210    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

2.  Decomposition  on  heating.    When  nitric  acid  is  boiled 
or  when  it  is  exposed  for  some  time  to  sunlight,  it  suffers 
a  partial  decomposition  according  to  the  equation 

4HNO3 — >-2H20  +  4N02  +  O2 

The  substance  NO2  (called  nitrogen  dioxide)  is  a  brownish 
gas  which  is  readily  soluble  in  water  and  in  nitric  acid.  It 
therefore  dissolves  in  the  undecomposed  acid,  and  imparts 
to  it  a  yellowish  or  reddish  color.  Concentrated  nitric  acid 
highly  charged  with  this  substance  is  called/MT/mt^  nitric  acid. 

3.  Oxidizing  action.    Because  of  its  easy  decomposition, 
nitric   acid  is   a  good  oxidizing   agent.     Under   ordinary 
circumstances,   when  acting   as   an   oxidizing   agent,  it  is 
decomposed  according  to  the  equation 

2  HNO3 >•  H20  +  2  NO  +  3  [O] 

The  oxygen  is  taken  up  by  the  substance  oxidized  and 
is  not  set  f ree^  which  fact  is  indicated  in  the  equation  by 
placing  the  symbol  for  oxygen  in  brackets.  Thus,  if  carbon 
is  oxidized  by  nitric  acid,  the  oxygen  combines  with  carbon, 
forming  carbon  dioxide  (CO2) : 

C  +  2[O] >-CO2 

4.  Action  upon  metals.   All  of  the  metals,  with  the  excep- 
tion of  gold,  platinum,  and  a  few  of  the  rare  metals,  are 
acted  upon  more  or  less  readily  by  nitric  acid.    The  action 
may  be  regarded  as  taking  place  in  two  different  ways, 
depending  upon  the  metal  and  the  concentration  of  the  acid. 

(a)  Metals  occurring  above  hydrogen  in  the  electrochemical 
series.  It  will  be  recalled  that  those  metals  occurring  above 
hydrogen  in  the  electrochemical  series  liberate  hydrogen 
from  dilute  acids.  Experiments  show,  however,  that  when 
these  metals  are  dissolved  in  nitric  acid,  hydrogen,  if  evolved 


COMPOUNDS  OF  NITROGEN  211 

at  all,  is  evolved  only  when  the  acid  is  very  dilute.  In  place 
of  hydrogen  certain  reduction  products  of  nitric  acid  are 
formed,  most  frequently  nitric  oxide,  a  gas  having  the 
formula  NO.  A  little  reflection  will  show  that  this  is  just 
what  we  might  expect,  for  hydrogen  is  a  strong  reducing 
agent,  while  nitric  acid  is  an  equally  strong  oxidizing  agent. 
Accordingly,  if  hydrogen  were  to  be  liberated  in  contact 
with  nitric  acid,  it  would  not  be  evolved  as  such,  but  would 
at  once  react  with  the  acid,  probably  in  the  following 


In  all  such  cases  the  complete  reaction  taking  place  may 
be  illustrated  by  the  following  equations  representing  the 
reaction  between  zinc  and  nitric  acid  (the  expression  2  [H] 
indicates  that  the  hydrogen  formed  is  not  liberated  as  such, 
but  as  fast  as  formed  reacts  according  to  equation  (2))  : 

Zn  +  2  HN03  -  >•  Zn(NO8)2  +  2  [H]  (1) 

3  [H]  +  HNO3  —  *-  2  H2O  +  NO  (2) 

In  case  the  nitric  acid  is  very  concentrated,  the  nitric 
oxide  is  oxidized  by  the  acid  to  form  nitrogen  dioxide,  — 
a  reddish-brown  gas  having  the  formula  NO2. 

It  is  often  convenient  to  express  in  a  single  equation  a 
reaction  that  really  takes  place  in  steps  such  as  that  between 
zinc  and  nitric  acid.  This  is  readily  done  by  combining  equa- 
tions (1)  and  (2)  as  given  above.  Before  the  equations  are  com- 
bined, however,  they  must  be  modified  so  as  to  express  the  fact 
that  all  the  hydrogen  represented  as  being  formed  according  to 
equation  (1)  reacts  with  the  nitric  acid  according  to  equation  (2). 
This  may  be  done  by  multiplying  the  first  equation  by  3  and 
the  second  equation  by  2.  The  two  equations  will  then  be  as 
follows  :  3  Zn  +  6  HN08  -  ^  3  Zn(N08)a  +  6  [H] 
6  [H]  +  2  HN08  -  *•  4  H20  +  2  NO 


212    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

By  canceling  the  common  factor  6  [H],  which  represents  the 
hydrogen  formed  in  the  one  reaction  and  used  up  in  the  other, 
and  then  combining  the  equations,  the  following  is  obtained : 

3  Zn  +  8  HN03 >•_  3  Zn(N03)2  +  4  H20  +  2  NO 

This  complete  equation  has  the  advantage  of  making  it  pos- 
sible to  calculate  very  easily  the  proportions  in  which  the 
various  substances  enter  into  the  reaction  or  are  formed  in  it. 
It  is  unsatisfactory  in  that  it  does  not  give  full  information 
about  the  way  in  which  the  reaction  takes  place. 

(b)  Metals  occurring  below  hydrogen  in  the  electrochemical 
series.  Those  metals  occurring  below  hydrogen  in  the 
electrochemical  series,  if  acted  upon  by  nitric  acid  at  all, 
are  first  oxidized  to  the  corresponding  oxides.  These  ox- 
ides, as  fast  as  formed,  react  with  more  acid  to  form  the 
corresponding  salt.  Thus,  with  copper  and  nitric  acid  the 
reactions  are  represented  by  the  following  equations : 

2  HNO3  — *  H2O  +  2  NO  +  3  [O]         (1) 

3[O]  +  3Cu — >-3CuO  (2) 

3  CuO  +  6  HN03  — *  3  Cu(NO3)2  +  3  H2O          (3) 

By  canceling  the  factors  3  [O]  and  3  CuO,  which  represent 
substances  formed  in  one  reaction  and  used  up  in  another, 
and  combining  the  three  equations  we  get  the  following: 

3  Cu  +  8  HN08  — +  3  Cu(NO3)2  +  2  NO  +  4  H2O 

Uses  of  nitric  acid.  Nitric  acid  is  used  in  very  large 
quantities  in  the  manufacture  of  explosives,  of  celluloid, 
and  of  dyes.  It  is  a  very  important  reagent  in  chemical 
laboratories. 

Aqua  regia.  Since  nitric  acid  is  a  good  oxidizing  agent, 
we  might  expect  it  to  liberate  chlorine  from  hydrogen 
chloride,  and  this  is  found  to  be  the  case.  A  mixture  of 


COMPOUNDS  OF  NITKOGEN  213 

1  part  of  nitric  acid  and  8  parts  of  hydrochloric  acid 
is  called  aqua  regia,  and  is  one  of  the  strongest  solvents 
known.  It  owes  its  solvent  powers  not  to  its  acid  proper- 
ties, but  to  the  nascent  chlorine  which  it  liberates.  Metals 
such  as  gold  and  platinum,  which  are  not  soluble  in  any 
of  the  common  acids,  readily  dissolve  in  aqua  regia,  being 
converted  into  chlorides  by  the  nascent  chlorine. 

Salts  of  nitric  acid;  nitrates.  The  salts  of  nitric  acid 
are  called  nitrates.  Many  of  these  salts  will  be  described 
in  the  study  of  the  metals.  They  are  all  soluble  in  water, 
and  when  heated  to  a  high  temperature,  undergo  decompo- 
sition. In  a  few  cases  a  nitrate,  on  being  heated,  evolves 
oxygen,  forming  a  nitrite: 

2  NaNO3 >•  2  NaNO2  +  O2 

In  most  cases  the  decomposition  goes  farther,  and  the 
metal  is  left  as  an  oxide : 

2  Pb(NO3)2  — *•  2  PbO  +  4  NO2  +  O2 

The  nitrates  are  especially  used  in  the  manufacture  of 
gunpowder,  sulfuric  acid,  nitric  acid,  and  as  a  fertilizer. 

Nitrous  acid  (HN02).  It  is  an  easy  matter  to  obtain  sodium 
nitrite  (NaNOj)  by  heating  sodium  nitrate,  as  explained  in  the 
previous  paragraph.  Now  when  sodium  nitrite  is  treated  with 
an  acid,  such  as  sulfuric  acid,  it  is  decomposed  and  nitrous 
acid  is  set  free : 

NaNOa  +  HaS04 >•  NaHS04  +  HNO2 

The  acid  is  very  unstable,  however,  and  decomposes  into  water 
and  oxides  of  nitrogen.  Sodium  nitrite  is  used  in  the  manu- 
facture of  coal-tar  dyes. 


214    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

OXIDES  OF  NITROGEN 

The  most  important  of  the  oxides  of  nitrogen  are  the 
following : 

Nitrous  oxide  (N2O)   ....  a  colorless  gas 

Nitric  oxide  (NO)  .    '.     .     .     .  a  colorless  gas 

Nitrogen  dioxide  (NO2)  ...  a  reddish-brown  gas 

Nitrogen  trioxide  (N2O3)      .     .  stable  only  at  low  temperatures 

Nitrogen  tetroxide  (N2O4)    .     .  a  low-boiling,  nearly  colorless  liquid 

Nitrogen  pentoxide  (N2O5)  .     .  a  white  solid 

Nitrous  oxide  (laughing  gas)  (N20).  Nitrous  oxide  was 
first  prepared  by  Priestley  in  1772.  Davy  determined  its 
composition  in  1800  and  was  the  first  to  point  out  the 
property  which  the  gas  possesses  of  rendering  one  tempo- 
rarily unconscious  when  it  is  inhaled. 

The  most  convenient  method  for  its  preparation  consists 
in  heating  ammonium  nitrate.  Just  as  ammonium  nitrite, 
when  heated,  yields  water  and  nitrogen  (p.  130),  so  am- 
monium nitrate  decomposes  in  a  similar  way,  forming 
water  and  nitrous  oxide.  The  similarity  between  the  two 
reactions  is  shown  in  the  following  equations: 

NH4N02— ^2H20  +  N2 
NH4NO3  — >•  2  H2O  +  N2O 

Nitrous  oxide  is  a  colorless  gas  somewhat  soluble  in 
water,  and  in  solution  has  a  slightly  sweetish  taste.  When 
inhaled  it  produces  a  kind  of  hysteria  (hence  the  name 
laughing  gas)  and,  if  taken  in  large  amounts,  insensibility 
to  pain,  and  unconsciousness.  It  was  the  first  substance  to 
be  used  as  an  anesthetic  in  surgery  and  is  still  used  in 
minor  operations,  such  as  those  of  dentistry. 

Nitrous  oxide  is  a  very  energetic  oxidizing  agent.  Sub- 
stances such  as  carbon,  sulfur,  iron,  and  phosphorus  burn 


COMPOUNDS  OF  NITROGEN 


215 


in  it  almost  as  brilliantly  as  in  oxygen,  forming  oxides 
and  setting  nitrogen  free.  Evidently  the  oxygen  in  nitrous 
oxide  is  not  held  in  very  firm  combination  by  the  nitrogen. 
Nitric  oxide.  Nitric  oxide  is  most  readily  prepared  by 
the  action  of  nitric  acid  (density,  1.2)  upon  certain  metals 
below  hydrogen  in  the  electrochemical  series,  such  as 
copper  (p.  212). 

The  inetui  is  placed  in  a  flask  A  (Fig.  88)  and  the  acid  slowly 
added  through  the  funnel  tube  B.  The  gas  escapes  through  C 
and  is  collected  over  water.  ._. 

The  gas  at  first  evolved 
combines  with  the  oxygen 
of  the  air  contained  in  the 
flask  to  form  the  reddish- 
brown  nitrogen  dioxide, 
but  this  is  dissolved  as  it 
bubbles  through  the  water. 


Nitric  oxide  is  a  color- 
less gas  slightly  heavier 
than  air.  It  is  a  much 
more  stable  compound 


FIG.  88.  The  preparation  of  nitric  oxide 


than  nitrous  oxide ;  nevertheless  it  can  be  decomposed  into 
its  elements  without  difficulty.  If  a  bit  of  phosphorus  is 
barely  ignited  and  at  once  introduced  into  a  jar  of  the  gas, 
the  flame  is  extinguished.  On  the  other  hand,  if  the  phos- 
phorus is  first  heated  until  vigorous  combustion  ensues,  and 
is  then  introduced  into  the  gas,  the  combustion  continues 
with  great  brilliancy. 

When  nitric  oxide  comes  into  contact  with  oxygen  or  air, 
it  at  once  combines  with  the  oxygen,  even  at  ordinary  tem- 
peratures, forming  a  reddish-brown  gas,  NO2,  which  is  called 
nitrogen  dioxide : 


2NO  +  C) 


2ND. 


216    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


To  show  the  formation  of  nitrogen  dioxide  from  nitric  oxide 
and  oxygen,  a  tube  is  filled  with  the  oxide,  inverted  in  water, 
and  pure  oxygen  is  passed  into  it,  as  shown  in  Fig.  89.  As 
each  bubble  of  oxygen  enters,  it  unites  with  the  nitric  oxide  to 
form  the  reddish-brown  dioxide.  In  a  few  minutes  the  color 
fades  (because  of  the  action  of  water  upon  the  dioxide),  and 
the  water  slowly  rises  in  the  tube. 

Nitrogen  dioxide  (N02).  This  gas,  as  we  have  just  seen,  is 
formed  by  allowing  nitric  oxide  to  come  into  contact  with 
oxygen.  It  can  also  be  made 
by  heating  certain  nitrates,  such 
as  lead  nitrate  (p.  213).  It  is 
a  reddish-brown  gas  of  unpleas- 
ant odor,  and  is  poisonous  when 
inhaled.  It  gives  up  a  part  of 
its  oxygen  to  burning  sub- 
stances, acting  as  an  oxidizing 
agent : 

N0 


FIG.  89.  The  formation  of  nitro-          Nitrogen    tetroxide.     At    lower 
gen  dioxide  from  nitric  oxide     temperatures     nitrogen     dioxide 
and  nitrogen  becomes  paler  in  color  and  con- 

denses   to  a   pale-yellow  liquid. 

It  has  been  shown  that  this  paler  gas  has  the  formula  N204, 
and  it  is  called  nitrogen  tetroxide.  At  ordinary  temperatures 
the  gas  is  a  mixture  of  the  two,  and  we  may  express  this 
relation  thus : 


Nitrogen  dioxide,  2NOa: 

high  temperatures 


:  nitrogen  tetroxide,  N204 
low  temperatures 


Acid  anhydrides.  The  oxides  N2O8  (nitrogen  trioxide) 
and  N2O5  (nitrogen  pentoxide)  are  rarely  prepared  and  need 
not  be  separately  described.  They  bear  a  very  interesting 


COMPOUNDS  OF  NITROGEN  217 

relation  to  the  acids  of  nitrogen.    When  dissolved  in  water 
they  combine  with  the  water  and  form  acids: 

N208 


Many  other  oxides  act  in  the  same  way,  combining  with 
water  to  form  an  acid.  Such  oxides  are  called  acid  an- 
hydrides. An  acid  anhydride  may  therefore  be  defined  as 
an  oxide  that  combines  with  water  to  form  an  acid. 

EXERCISES 

1.  Perfectly  dry  ammonia  does  not  affect  litmus  paper.    Explain. 

2.  Can  ammonia  be  dried  by  passing  the  gas  through  concen- 
trated sulfuric  acid?    Explain. 

3.  Ammonium  hydroxide  is  a  weak  base;   that  is,  it  is  not 
highly  dissociated.     When  it  is  neutralized  by  strong   acids,  the 
heat  of  reaction  is  less  than  when  strong  bases  are  so  neutralized. 
Suggest  some  possible  cause  for  this. 

4.  Write   the   equations  for  the  reactions  taking  place  when 
ammonium  hydroxide  is  neutralized  by  hydrochloric  acid,  by  sul- 
furic acid,  and  by  nitric  acid  respectively. 

5.  It  is    said   that    nitric    acid   is  formed   in   the   air   during 
thunderstorms.    How  would  you  account  for  its  formation? 

6.  W hat  does  the  word  ammonia  mean  ?    (Consult  dictionary.) 

7.  Why  is  nitric  acid  said  to  be  a  strong  acid? 

8.  What  are  the  properties  of  ammonia  that  make  it  suitable 
for  use  in  the  preparation  of  artificial  ice? 

9.  Write  the  equations  representing  the  reactions  between  am- 
monium   hydroxide    and    sulfuric    acid    and    between    ammonium 
hydroxide  and  nitric  acid,  in  accordance  with  the  theory  of  electro- 
lytic dissociation. 

10.  State  the  compounds  and  ions  present  in  aqua  ammonia. 

11.  What  is  meant \>y  the  statement  "the  reaction  between  water 
and  ammonia  is  a  reversible  reaction"? 


218    AN  ELEMENTARY  STUDY  OF  CHEMISTEY 

12.  Why  is  it  necessary  to  apply  heat  in  the  preparation  of  nitric 
acid  from  sodium  nitrate  ? 

13.  Give  the  steps  in  the  production  of  nitric  acid  from  air  and 
water. 

14.  How  many  liters  of  ammonia  at  0°  and  760  mm.  pressure 
will  1 1.  of  water  dissolve  ?    What  would  this  volume  of  ammonia 
weigh  ?   What  weight  of  ammonium  chloride  would  be  necessary  to 
prepare  it? 

15.  (a)  Calculate   the  weight  of  1  1.  of  the  concentrated   nitric 
acid  of  commerce  (p.  209).     (ft)  What  weight  of  hydrogen  nitrate 
(HNO3)  will  this  contain  ?    (c)  What  weights  of  materials  (H2SO4 
and  NaNO3)  are  necessary  for  its  preparation? 

16.  How  many  liters  of  nitrous  oxide,  measured  under  standard 
conditions,  can  be  prepared  from  10  g.  of  ammonium  nitrate? 

17.  What  weight  of  copper  is  necessary  to  prepare  50  1.  of  nitric 
oxide? 

18.  What  weight  of  sodium  hydroxide  is  necessary  to  neutralize 
1  1.  of  the  concentrated  nitric   acid  of  commerce?    Calculate  the 
weight  of  the  resulting  compound. 

19.  40  cc.  of  the  concentrated  nitric  acid  of  commerce  was  added 
to  25  g.  of  zinc.    After  the  reaction  had  ceased,  the  resulting  mass 
was  heated  until  the  water  present  was  all  evaporated.    Calculate 
the  weight  of  the  residue. 


CHAPTER  XIX 
SPEED  OF  REACTIONS;    EQUILIBRIUM 

Speed  of  a  reaction.  If  the  reactions  we  have  studied 
are  attentively  examined,  it  will  be  noticed  that  some  go 
on  very  slowly,  while  others  are  very  rapid.  For  example, 
Lavoisier  found  that  tin  must  be  heated  in  the  air  for 
days  before  an  appreciable  quantity  of  the  oxide  is  formed 
(p.  4).  So,  too,  the  action  of  acids  upon  pure  zinc  is  very 
slow.  On  the  other  hand,  when  iron  and  sulfur  have  been 
heated  to  the  temperature  at  which  noticeable  action  be- 
gins, the  action  proceeds  very  rapidly.  We  express  these 
facts  by  saying  that  reactions  differ  greatly  among  them- 
selves in  their  speed.  By  the  speed  of  a  reaction  is  meant 
the  quantity  of  a  given  substance  that  undergoes  change 
in  a  unit  of  time. 

Factors  that  affect  the  speed  of  a  reaction.  A  number 
of  factors  alter  the  speed  at  which  a  given  reaction  will 
take  place. 

1.  Temperature.  The  most  familiar  way  to  hasten  a  reac- 
tion is  to  raise  the  temperature  of  the  materials  under- 
going change.  In  ordinary  combustion  it  is  a  familiar  fact 
that  the  hotter  the  fire,  the  faster  the  coal  burns,  while  at 
ordinary  temperatures  the  coal  does  not  seem  to  oxidize 
at  all,  though  there  is  plenty  of  oxygen  in  contact  with 
it.  It  is  probable  that  even  under  these  latter  conditions 
there  is  some  chemical  action,  but  it  is  so  slow  that  we 
cannot  easily  detect  it.  In  general,  it  has  been  found  that 

219 


220    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

a  rise  of  10°  about  doubles  the  speed  of  most  reactions. 
Consequently,  if  we  are  in  doubt  as  to  whether  two  things 
can  act  upon  each  other,  we  gradually  raise  the  tempera- 
ture of  the  mixture,  knowing  that  this  will  increase  the 
speed  of  action  if  there  is  any  taking  place. 

2.  Concentration  —  mass  action.  It  is  evident  that  if  we 
increase  the  quantity  of  oxygen  in  contact  with  the  sur- 
face of  a  burning  substance,  we  shall  increase  the  speed  of 
the  combustion.  For  this  reason  substances  burn  much 
more  rapidly  in  pure  oxygen  than  in  air,  for  the  latter  is 
only  one  fifth  oxygen.  The  blacksmith  pumps  air  into 
the  forge  to  supply  oxygen  faster  than  it  would  be  sup- 
plied by  natural  draft,  and  this  operation  amounts  to 
increasing  the  quantity  of  oxygen  available  at  any  one 
time.  We  can  express  these  facts  by  saying  that  the  speed 
with  which  a  given  material  burns  depends  upon  the  con- 
centration of  the  oxygen.  By  concentration  we  mean  the 
weight  of  oxygen  present  in  a  given  volume  of  gas. 

Instead  of  increasing  the  concentration  of  the  oxygen, 
we  may  increase  the  surface  of  the  burning  substance.  A 
log  of  wood  burns  more  slowly  than  the  same  log  split  into 
kindling  or  worked  into  shavings.  A  lump  of  coal  burns 
rather  slowly;  but  when  finely  powdered  and  suspended 
in  the  air  as  dust,  it  burns  almost  instantaneously  and 
with  explosive  violence.  If  we  burn  a  given  weight  of 
coal  in  the  air,  we  shall,  of  course,  get  a  definite  weight 
of  product  and  use  up  a  definite  weight  of  oxygen.  But 
if  we  force  into  contact  with  the  coal  much  more  oxygen 
than  can  be  used,  the  coal  burns  much  faster.  This  effect 
of  the  excess  of  oxygen  is  called  mass  action.  In  general, 
any  reaction  can  be  hastened  by  the  effect  of  mass  action ; 
that  is,  by  using  a  large  excess  of  one  of  the  reacting 
substances. 


SPEED  OF  REACTIONS;  EQUILIBRIUM       221 

3.  Catalysis.  We  have  found  that  certain  reactions  go 
on  much  more  rapidly  in  the  presence  of  some  material 
that  does  not  seem  to  take  any  real  part  in  the  process, 
since  it  remains  left  over  at  the  end.  Such  a  substance  is 
called  a  catalyzer,  and  its  action  is  called  catalysis  (p.  26). 
Thus,  at  a  given  temperature  potassium  chlorate  gives  up 
oxygen  more  rapidly  when  manganese  dioxide  has  been 
added  to  it  than  when  the  pure  substance  is  heated  (p.  26). 
Nitrogen  combines  with  hydrogen  much  more  rapidly  iii 
the  presence  of  iron  than  in  its  absence  (p.  203).  A  cata- 
lyzer reminds  us  of  a  lubricant  in  machinery.  It  makes 
the  reaction  proceed  more  rapidly,  just  as  oil  makes  the 
machine  run  more  easily. 

Reversible  reactions.  We  have  met  with  a  number  of 
reactions  that  are  especially  interesting,  because  by  changing 
the  conditions  we  can  make  them  go  in  either  direction  at 
will.  Thus,  when  we  heat  mercuric  oxide  we  obtain  mer- 
cury and  oxygen  (p.  25),  while  if  we  heat  mercury  in 
contact  with  oxygen  at  a  somewhat  different  temperature, 
we  obtain  mercuric  oxide.  These  facts  are  represented  in 
the  following  way: 


In  a  similar  way  we  have  found  that  when  an  electric 
discharge  is  passed  through  a  mixture  of  nitrogen  and 
hydrogen  we  get  a  small  quantity  of  ammonia,  yet  when 
the  discharge  is  passed  through  ammonia  we  get  a  mixture 
of  nitrogen  and  hydrogen  (p.  204)  : 


When  ammonia  is  dissolved  in  water  we  have  every  reason 
for  thinking  that  ammonium  hydroxide  is  formed  (p.  204), 


222    AN  ELEMENTARY  STUDY  OF  CHEMISTEY 

yet  when  we  attempt  to  evaporate  the  solution  to  obtain 
this  compound,  it  decomposes  into  water  and  ammonia  : 


Reactions  of  this  kind  are  called  reversible  reactions. 

Equilibrium.  If  we  remember  that  the  materials  taking 
part  in  a  reaction  are  made  up  of  a  vast  number  of  mole- 
cules, all  of  which  are  in  motion  at  various  speeds  and 
are  constantly  changing  then-  relations  to  each  other,  it 
will  not  be  difficult  to  imagine  why  some  molecules  in 
a  given  mixture  may  be  decomposing  while  others  are 
being  formed.  Moreover,  at  the  outset  of  the  reaction 
the  masses  of  the  substances  present  tend  to  drive  the 
reaction  in  one  direction,  while  later,  when  the  masses 
of  the  original  substances  have  diminished,  the  masses  of 
the  substances  formed  have  increased.  If  these  latter  are 
capable  of  reacting  to  form  the  original  substances,  the 
effect  of  mass  will  now  be  to  drive  the  reaction  in  the 
reverse  direction.  In  time  a  condition  will  be  reached 
in  which  the  changes  taking  place  in  the  one  direction 
will  just  offset  those  in  the  other.  The  average  percentage 
of  each  material  present  will  then  remain  unchanged, 
though  the  individual  molecules  will  keep  on  changing. 
This  condition  of  affairs  is  called  equilibrium.  Thus, 
ammonia,  hydrogen,  and  nitrogen  come  to  equilibrium 
in  the  presence  of  electric  discharge  when  there  is  about 
7  per  cent  of  ammonia  present,  the  percentage  depending 
upon  the  temperature  and  the  relative  quantities  of  the 
gases  originally  mixed. 

Effect  of  mass  upon  equilibrium.  When  equilibrium  has 
been  reached,  suppose  we  add  an  additional  quantity  of 
one  of  the  acting  substances  —  say,  hydrogen  in  the  case 
just  mentionc-1  This  will  enable  the  nitrogen  to  act  more 


SPEED  OF  REACTIONS;  EQUILIBRIUM       223 

rapidly  upon  the  hydrogen,  for  the  two  kinds  of  molecules 
will  now  meet  more  frequently.  It  will  not  at  all  affect 
the  rate  at  which  ammonia  is  decomposing,  for  this  does  not 
in  any  way  depend  upon  the  presence  of  hydrogen.  The 
net  effect  will  therefore  be  to  bring  about  a  new  equilibrium 
in  which  a  larger  percentage  of  ammonia  is  present. 

Changing  an  equilibrium  to  a  completed  reaction.  If  we 
were  to  withdraw  the  ammonia  as  fast  as  it  is  formed, 
before  it  has  time  to  decompose,  the  reaction  ought  to  go 
on  until  either  the  hydrogen  or  the  nitrogen  is  used  up. 
The  ammonia  can  be  so  withdrawn  by  inclosing  the  gases 
over  water  containing  acid  during  the  discharge,  and  in 
this  case  the  reaction  goes  on  until  one  or  the  other  react- 
ing gases  is  used  up.  The  point  of  equilibrium  can  there- 
fore be  altered  or  the  equilibrium  changed  into  a  completed 
reaction,  by  changing  the  relative  masses  of  the  substances 
taking  part  in  the  reaction. 

Equilibrium  in  solution.  In  aqueous  solution  we  are 
interested  chiefly  in  the  equilibrium  of  ions.  The  mole- 
cules of  an  electrolyte  keep  dissociating  into  ions,  while 
the  ions,  on  meeting,  recombine  to  form  molecules,  the 
result  being  an  equilibrium  between  the  two  conditions. 
Thus,  with  nitric  acid  we  have  the  equilibrium 


If  we  mix  two  electrolytes,  the  equilibrium  that  is  reached 
is  a  much  more  complicated  one,  for  any  positive  ion  may 
unite  with  any  negative  one.  At  equilibrium  all  possible 
ions  and  combinations  of  ions  will  be  present.  Thus,  when 
we  mix  sodium  nitrate  and  sulfuric  acid  in  the  preparation 
of  nitric  acid,  we  have  present  the  ions  Na+,  NOg~,  H+, 
and  SO4~-,  together  with  the  molecules  NaNO8,  Na2SO4, 
NaHSO4,  HNO8,  and  H2SO4. 


224    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Completion  of  reactions  in  solution.  The  chemist  makes 
use  of  reactions  to  secure  various  compounds  in  pure  con- 
dition, and  he  wishes  the  yield  to  be  as  large  as  possible. 
Reactions  which  stop  short  of  completion  and  end  in  an 
equilibrium  are  not  suited  to  manufacturing  purposes 
unless  means  can  be  found  to  break  up  the  condition  of 
equilibrium  and  bring  the  reaction  to  a  definite  conclusion. 
There  are,  in  general,  three 
conditions  under  which  this 
may  be  accomplished. 

1.  A  volatile  gas  may  be 
formed.  If  the  reaction  is  con- 
ducted under  conditions  such 
that  one  of  the  products  is 
a  gas  insoluble  in  the  solvent, 
the  gas  will  make  its  escape 
as  fast  as  it  is  formed.  The 
reaction  will  then  continue 
until  one  or  the  other  of  the 
ions  taking  part  in  the  forma- 
tion of  the  gaseous  molecules 

'IS  use(^  UP' 

For  example,  when  we  mix 

sulfuric  acid  and  sodium  ni- 

trate, no  visible  reaction  takes  place.  But  if  we  heat 
the  mixture  above  the  boiling  point  of  nitric  acid,  all  of 
this  substance  formed  in  the  equilibrium  between  the  two 
ions  H+  and  NO3~  is  converted  into  a  gas,  insoluble  in 
sulfuric  acid.  The  nitric  acid  distills  away  until  all  of 
the  NOg~  ions  are  used  up.  We  then  have  a  completed 
reaction  expressed  in  the  equation 


FIG.  90.    The    precipitation   of 
silver  chloride 


NaHSO 


SPEED  OF  REACTIONS;  EQUILIBRIUM       225 

2.  An  insoluble  solid  may  be  formed.  When  hydrogen 
chloride  (HC1)  and  silver  nitrate  (AgNO8)  are  brought 
together  in  solution,  we  have  the  two  kinds  of  molecules 
just  named,  the  ions  H+,  Cl~,  Ag+,  NO3~,  and  the  new 
combinations  HNO3  and  AgCl.  One  of  these,  namely 
silver  chloride  (AgCl),  is  insoluble  in  water,  and  as  fast 
as  it  is  formed  it  separates  from  the  solution  as  a  curdy 
white  precipitate  (Fig.  90).  The  reaction  therefore  con- 
tinues until  either  the  ion  Ag+  or  the  ion  Cl~  is  used  up, 
the  completed  equation  being 


+  (Ag+  +  NO,-)  —  >•  H+  +  N03-  +  AgCl 

3.  Two  different  ions  may  unite  to  form  an  undissociated 
molecule.  When  we  bring  together  in  solution  sodium 
hydroxide  and  hydrochloric  acid,  we  have  the  ions  H+> 
C1-,  Na+,  and  OH~.  The  H+  ions  and  the  OH~  ions 
unite  to  form  molecules  of  water  which  do  not  again  part 
into  ions  save  to  a  very  slight  extent.  This  leaves  only 
the  ions  of  NaCl  in  solution,  the  equation  being 


(Na+  +  OH-)  +  (H+  +  C1-)  —  *  H2O  +  Na 

Neutralization  is  practically  a  completed  reaction  because 
water  is  so  little  ionized. 

The  preparation  of  acids.  The  principle  explained  in 
(1)  is  very  frequently  applied  in  the  preparation  of  various 
acids.  Most  of  the  substances  that  form  acids  in  solution 
have  rather  low  boiling  points,  while  concentrated  sulfuric 
acid  has  a  rather  high  boiling  point  (338°).  Consequently, 
if  we  take  a  salt  of  almost  any  common  acid  —  a  nitrate, 
a  chloride,  or  an  acetate  —  and  treat  it  with  concentrated 
sulfuric  acid,  at  the  same  time  heating  the  solution,  the  vola- 
tile acid  will  boil  out  of  the  mixture  and  can  be  condensed 
by  cooling  the  vapor  as  in  the  preparation  of  nitric  acid. 


226    AN  ELEMENTAKY  STUDY  OF  CHEMISTRY 

Hydrolysis.  While  water  is  very  little  ionized,  neverthe- 
less it  forms  some  ions.  Moreover,  when  a  salt  is  dissolved 
in  water  to  form  a  dilute  solution,  the  relative  mass  of 
the  water  is  very  great.  The  reaction  of  neutralization  is 
therefore  reversed  to  a  slight  extent,  forming  a  small  amount 
of  free  base  and  of  free  acid,  thus: 

NaNO3  +  H2O  +^  NaOH  +  HNO2 

A  reaction  of  this  kind,  in  which  water  acts  upon  a  salt 
to  form  a  base  and  an  acid,  is  called  hydrolysis.  If  the 
base  formed  in  hydrolysis  is  very  weak  and  the  acid  is 
strong,  the  solution  will  turn  blue  litmus  red,  as  is  true 
with  the  salts  of  aluminium.  If  the  base  is  very  strong 
and  the  acid  weak,  the  solution  will  turn  red  litmus  blue, 
as  is  the  case  with  many  salts  of  sodium.  If  both  the 
acid  and  the  base  formed  are  weak,  then  the  compound 
may  be  completely  hydrolyzed. 

EXERCISES 

1.  Can  you  mention  any  reversible  reactions,  other  than  those 
given  in  this  chapter? 

2.  Suggest  a  method  for  the  preparation  of  hydrogen  chloride. 

3.  Would  silver  nitrate  produce  a  precipitate  when  added  to  a 
solution  of  sodium  chloride  (common  salt)  ?    If  so,  how  would  the 
precipitate  compare  in  composition  with  that  produced  when  silver 
nitrate  is  added  to  hydrochloric  acid  ? 

4.  Barium  sulfate  (BaSO4)  is  a  white  insoluble  compound  much 
used  as  a  pigment  in  making  paints.  Suggest  a  method  for  preparing  it. 

5.  Is  the  reaction  NH3  +  H2O »-NH4OH  reversible?    If  so, 

state  the  conditions  under  which  it  will  go  in  each  direction. 

6.  Is  the  reaction  expressed  by  the  equation  2  H2  +  O2 >-2  H2O 

reversible  ?   If  so,  state  the  conditions  under  which  it  will  go  in  each 
direction. 


SPEED  OF  EEACTIONS;  EQUILIBRIUM       227 

7.  Carbonic  acid  is  a  very  weak  acid,  while  sodium  hydroxide  is 
a  strong  base.    How  will  a  solution  of  sodium  carbonate  act  towards 
litmus  paper  ? 

8.  (a)  In  an  experiment  10  g.  of  iron  filings  and  8  g.  of  sulfur 
were  mixed  and  heated  until  reaction  started.    What  weight  of 
iron  sulfide  was  formed  in  the  reaction  ?    (ft)  In  a  second  experiment 
10  g.  of  iron  and  16  g.  of  sulfur  were  heated  in  a  similar  way.    What 
weight  of  iron  sulfide  was  formed  ?    (e)  What  difference  would  you 
expect  to  find  in  the  two  experiments  ? 

9.  In  the  preparation  of  hydrogen  chloride  from  sodium  chloride 
and  sulfuric  acid,  why  does  the  reaction  go  to  practical  completion  ? 


CHAPTER  XX 
SULFUR;   SELENIUM;   TELLURIUM 

History  and  occurrence.  Sulfur  occupied  a  prominent 
place  among  the  few  elements  known  to  the  ancients,  and 
played  an  important  part  in  the  older  views  concerning 
the  composition  of  matter.  It  occurs  in  nature  in  both 
the  free  and  the  combined  condition. 

In  certain  volcanic  regions,  especially  in  Sicily,  large 
deposits  of  free  sulfur  are  found,  which  until  recent  times 


FIG.  91.   Forcing  liquid  sulfur  from  deep  wells  in  Louisiana  by  means 
of  compressed  air 

constituted  the  principal  source  of  the  world's  supply  of 
this  element.  Free  sulfur  also  occurs  in  Japan,  Spain, 
Iceland,  Mexico,  and  in  different  localities  in  the  United 
States,  especially  in  Louisiana.  In  combination,  sulfur 
occurs  abundantly  in  the  form  of  sulfides  and  sulfates.  In 
smaller  amounts  it  is  found  in  a  great  variety  of  minerals 
and  is  a  constituent  of  many  vegetable  and  animal  sub- 
stances, especially  of  the  yolk  of  eggs. 


SULFUR;  SELENIUM;  TELLURIUM 


229 


Extraction  of  sulfur.  In  Louisiana  the  sulfur  occurs  in 
deposits  far  underground  and  covered  with  quicksand  so 
that  it  cannot  be  mined.  One  of  these  deposits  lies  at  a 
depth  of  700  feet,  is  circular  in  shape,  and  is  about  half 
a  mile  in  diameter  and  500  feet  in  thickness.  Wells  are 
drilled  into  the  deposit,  and  superheated  water  is  forced 
down  through  suitable  pipes.  The  hot  water  melts  the 
sulfur,  which  is  then  forced  up  a  separate  pipe  by  com- 
pressed air  (A,  Fig.  91).  The  liquid  sulfur  then  solidifies 
in  very  large  blocks.  A  single  well  has  produced  500  tons 
daily,  and  the  product  is 
99.5  per  cent  pure.  About 
250,000  tons  are  produced 
annually  from  this  deposit. 

In  Sicily  a  very  simple 
but  wasteful  method  is  used 
to  separate  sulfur  from  the 


FIG.  92.    A  sulfur  still 


rock  and  earthy   materials 
with    which    it    is    mixed. 

The  material  is  piled  up  in  heaps  and  set  on  fire,  and  the 
heat  from  the  burning  of  a  part  of  the  sulfur  melts  another 
portion,  which  collects  as  a  liquid  at  the  bottom  of  the  pile. 
This  is  drained  off  and  purified  by  distillation  in  a  retort 
(J()  (Fig.  92),  the  exit  tube  of  which  opens  into  a  cool- 
ing chamber  (J5)  of  brickwork.  When  the  sulfur  vapor 
first  enters  the  cold  chamber  it  condenses  as  a  fine  crys- 
talline powder  called  flowers  of  sulfur.  As  the  condensing 
chamber  becomes  warm  the  sulfur  condenses  as  a  liquid 
and  is  drawn  off  into  cylindrical  molds,  the  product  being 
called  roll  sulfur,  or  brimstone. 

Properties.  Sulfur  is  a  pale  yellow  solid  without  marked 
taste  and  with  but  a  faint  odor.  It  is  insoluble  in  water. 
It  melts  when  heated,  forming  a  thin,  straw-colored  liquid. 


230    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


If  the  temperature  is  gradually  raised,  this  liquid  turns 
darker  in  color  and  becomes  thicker  until,  at  about  235°,  it 
is  almost  black  and  is  so  viscous  that  a  vessel  containing  it 
can  be  inverted  without  danger  of  the  liquid's  flowing  out. 
At  higher  temperatures  it  becomes  mobile  again,  and  at 
444.6°  boils,  forming  a  yellowish  vapor.  When  the  hot  sul- 
fur is  cooled  the  same  changes  take  place  in  reverse  order. 
Varieties  of  sulfur.  Sulfur  exists  in  a  number  of  allo- 
tropic  forms  (p.  112),  which  are  easy  to  obtain.  The  best- 
known  are  the  following: 
1.  Ordinary,  or  rhombic, 
sulfur.  When  sulfur  crys- 
tallizes from  solution  in 
liquids  (notably  from  car- 
bon disulfide)  it  is  ob- 
tained in  compact  yellow 
crystals  which  melt  at 
112.8°  and  have  a  density 
of  2.06.  These  crystals 
usually  have  8  sides,  be- 
long to  the  rhombic  sys- 
tem, and  are  known  as 
rhombic  sulfur.  Sulfur  is  often  found  in  a  crystalline  state 
in  nature  (Fig.  93),  and  these  crystals  always  belong  to 
the  rhombic  system  (see  Appendix).  Brimstone  is  com- 
posed largely  of  rhombic  sulfur. 

2.  Prismatic,  or  monoclinic,  sulfur.  When  melted  sulfur  is 
allowed  to  cool  until  a  part  of  the  liquid  has  solidified,  and 
the  remaining  liquid  is  then  poured  off,  it  is  found  that  the 
solid  sulfur  remaining  in  the  vessel  is  in  the  form  of  fine 
needle-shaped  crystals,  which  melt  at  119.2°  and  have  a 
density  of  1.96.  The  needle-shaped  form  is  called  monoclinic 
sulfur,  since  the  crystals  belong  to  the  monoclinic  system. 


FIG.  93.    Crystals  of  rhombic  sulfur 
as  they  are  found  in  nature 


SULFUK;  SELENIUM;  TELLURIUM  231 

Relation  of  rhombic  to  monoclinic  sulfur.  Experiments 
have  shown  that  whenever  sulfur  crystallizes  at  ordinary 
temperature  the  rhombic  form  is  obtained;  when  crystal- 
lized at  higher  temperatures,  as  when  the  sulfur  is  melted 
and  allowed  to  cool,  the  monoclinic  form  is  obtained. 
Moreover,  the  temperature  below  which  sulfur  assumes 
the  rhombic  form  and  above  which  it  assumes  the  mono- 
clinic  form  is  a  perfectly  definite  one ;  namely,  95.5°.  At 
this  temperature,  known  as  the  transition  temperature,  the 
two  forms  of  crystals  remain  unchanged  when  in  contact 
with  each  other.  If  heated  above  95.5°,  the  rhombic  form 
gradually  changes  into  the  monoclinic  form;  if  cooled 
below  95.5°,  the  monoclinic  gradually  changes  into  the 
rhombic  form.  This  change  of  one  form  into  the  other 
ordinarily  takes  place  very  slowly,  so  that  some  days  may 
be  required  before  the  change  is  complete.  It  has  also 
been  found  that  an  increase  of  pressure  promotes  the 
change  from  rhombic  sulfur  into  monoclinic  sulfur. 

Amorphous  sulfur.  In  discussing  the  physical  properties 
of  sulfur,  attention  was  called  to  the  fact  that  sulfur 
is  easily  melted  and  forms  a  pale-yellow,  mobile  liquid, 
which  at  a  higher  temperature  becomes  dark  and  viscous. 
At  intermediate  temperatures  the  liquid  obtained  consists  of 
varying  amounts  of  the  mobile  liquid  (S\)  and  the  viscous 
liquid  (S/x)  in  equilibrium  with  each  other.  If  the  molten 
sulfur  is  heated  to  boiling  and  poured  into  cold  water, 
the  sudden  chilling  prevents  the  crystallization  of  the 
viscous  liquid,  so  that  an  amorphous,  dough-like  product 
is  obtained.  This  form  is  insoluble  in  carbon  disulfide 
and  is  known  as  plastic  sulfur.  It  is  simply  a  mixture  of 
the  two  liquid  forms  in  a  very  much  undercooled  state 
(p.  109).  On  standing,  plastic  sulfur  changes  in  part 
into  rhombic  crystals. 


232    AK  ELEMENTAEY  STUDY  OF  CHEMISTRY 


The  formation  of  plastic  sulfur  is  shown  in  a  very  striking 
manner  by  distilling  sulfur  from  a  small,  short-necked  retort 
(Fig.  94)  and  allowing  the  distillate  to  run  into  cold  water. 

Chemical  conduct  of  sulfur.  When  sulfur  is  heated  to 
ignition  in  oxygen  or  in  the  air,  it  burns  with  a  pale-blue 
flame  and  forms  sulfur  dioxide  (SO2).  Small  quantities  of 
sulfur  trioxide  (SO3)  may  also  be  formed  in  the  com- 
bustion of  sulfur.  Most  metals  when  heated  with  sulfur 
combine  directly  with  it,  forming 
metallic  sulndes.  In  some  cases 
the  action  is  so  energetic  that  the 
mass  becomes  incandescent,  as 
has  been  seen  in  the  case  of  iron 
uniting  with  sulfur.  This  prop- 
erty recalls  the  action  of  oxygen 
upon  metals,  and  in  general  the 
metals  which  combine  readily 
with  oxygen  are  apt  to  com- 
bine quite  readily  with  sulfur. 

Uses  of  sulfur.    Large  quan- 
tities of  sulfur  are  used  in  the 


FIG.  94.    The  preparation  of 
plastic  sulfur 


manufacture  of  gunpowder,  vulcanized  rubber,  carbon  di- 
sulfide,  sulfur  dioxide,  sulfuric  acid,  and  salts  of  various 
kinds.  It  is  also  used  extensively  in  the  manufacture  of 
insecticides  for  use  in  orchards  and  vineyards. 

Lime-sulfur  spray.  The  chief  sulfur  insecticide  is  known  as 
lime-sulfur  spray.  It  is  made  by  boiling  sulfur  with  slaked 
lime,  by  which  process  a  deep-red  liquor  is  obtained  which  con- 
sists essentially  of  a  solution  of  sulndes  of  calcium  (CaS4  and 
CaSR).  The  liquid  is  a  very  efficient  insecticide,  particularly 
for  scale,  and  it  is  also  a  fungicide.  Large  quantities  of  it  are 
used  for  spraying  fruit  trees. 


§ULFUR;  SELENIUM;  TELLURIUM 


233 


COMPOUNDS  OF  SULFUR  WITH  HYDROGEN 

The  following  compounds  of  sulfur  with  hydrogen  are 
known :  hydrogen  sulfide  (H2S),  a  foul-smelling  gas ; 
hydrogen  persulfide,  a  liquid  which  is  probably  a  mixture 
of  the  sulfides  H2S4  and  H2S6. 

Hydrogen  sulfide  (H2S).  Hydrogen  sulfide  is  present  in 
the  vapors  issuing  from  volcanoes.  Dissolved  in  water,  it 
constitutes  the  so-called  sulfur  waters  of  common  occur- 
rence. It  is  formed  when  organic  matter  containing  sulfur 
undergoes  decay,  and  the  disagree- 
able odor  attending  such  changes 
is  due  in  part  to  the  presence  of 
this  gas. 

Preparation.  Hydrogen  sulfide 
is  prepared  in  the  laboratory  by 
treating  a  sulfide  with  an  acid. 
Iron  sulfide  (FeS)  and  hydro- 
chloric acid  are  usually  employed : 


FeS  +  2HCl 


FeCl 


FIG.  95.   The  preparation  of 
hydrogen  sulfide 


A  convenient  apparatus  is  shown 
in   Fig.  95.    A  few  lumps  of  iron 

sultide  are  placed  in  the  bottle  A,  and  dilute  acid  is  added 
a  little  at  a  time  through  the  funnel  tube  B.  The  gas  escapes 
through  the  tube  C  and  may  be  collected  by  displacement  of 
air;  or  it  may  be  passed  into  water,  forming  a  solution.  The 
Kipp  generator  (Fig.  17)  is  more  convenient  than  the  above 
apparatus  if  a  larger  quantity  of  the  gas  is  desired. 

Properties.  Hydrogen  sulfide  is  a  colorless  gas  having 
a  mild,  disagreeable  taste  and  an  offensive  odor.  It  is 
1.18  times  as  heavy  as  air.  When  liquefied  it  boils  at 
-  61.6°  and  freezes  at  -  82°.  One  volume  of  water  at  15° 


234    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

dissolves  3.05  volumes  of  the  gas.  When  this  solution  is 
heated  to  boiling,  the  gas  is  all  expelled.  In  pure  form 
it  acts  as  a  violent  poison  and,  even  when  diluted  largely 
with  air,  produces  headache,  dizziness,  and  nausea.  For- 
tunately its  extremely  disagreeable  odor  gives  warning  of 
its  presence. 

Chemical  conduct.  1.  Acid  properties.  In  aqueous  solu 
tion  hydrogen  sulfide  is  slightly  dissociated,  giving  hydro- 
gen ions.  The  solution  therefore  acts  as  a  weak  acid  and 
is  known  as  hydro sulf uric  acid.  It  possesses  the  general 
properties  of  an  acid,  turning  blue  litmus  red  and  neu- 
tralizing bases  with  the  formation  of  sulfides. 

2.  Action  of  heat.    When  heated  to  a  high  temperature 
hydrogen  sulfide  is  decomposed  into  its  elements,  the  speed 
of  decomposition  being  marked  at  500°. 

3.  Action  of  oxygen.    When  a  solution  of  hydrogen  sul- 
fide in  water  is  exposed  to  the  air,  the  hydrogen  of  the 
sulfide  unites  with  oxygen  to  form  water,  while  the  sulfur 
is  liberated  and  settles  to  the  bottom  of  the  liquid.     In 
this  way  are  formed  the  deposits  of  sulfur  found  about 
many  springs.     At  a  high  temperature  hydrogen  sulfide 
burns  readily  in  either  oxygen   or  air  according  to  the 
equation         2  H2S  +  3  O2 +  2  H2O  +  2  SO2 

When  there  is  not  sufficient  oxygen  to  combine  with  both 
the  sulfur  and  the  hydrogen,  the  latter  element  combines 
with  the  oxygen  and  the  sulfur  is  set  free : 

2H2S  +  02 — ^2H20  +  2S 

4.  Reducing  action.     Because  of  the  hydrogen  present, 
together  with  the  ease  with  which  it  is  given  up  in  con- 
tact with  an  oxidizing  agent,  hydrogen  sulfide  acts  as  a 
strong  reducing  agent.    Thus,  when  it  is  bubbled  through 


SULFUR;  SELENIUM;  TELLURIUM  235 

concentrated  nitric  or  sulfuric  acid,  both  of  which  are 
strong  oxidizing  agents,  the  hydrogen  of  the  sulfide  com- 
bines with  a  portion  of  the  oxygen  of  the  acid  to  form 
water,  the  acid  being  at  the  same  time  reduced. 

A  much-used  method  of  drying  gases  consists  in  bubbling 
them  through  concentrated  sulfuric  acid,  which  absorbs  the 
moisture.  It  is  evident,  however,  from  the  statements  just 
made,  that  this  method  cannot  be  used  for  drying  hydrogen 
sulfide. 

5.  Action  on  metals.  Hydrogen  sulfide  acts  upon  many 
metals,  forming  sulfides.  Silver  sulfide  (Ag2S)  is  black,  and 
it  is  owing  to  traces  of  hydrogen  sulfide  in  the  air  that 
silver  objects  tarnish. 

Salts  of  hydrosulfuric  acid  ;  sulfides.  The  salts  of  hydro- 
sulfuric  acid,  or  sulfides,  form  an  important  class  of  com- 
pounds, and  many  occur  in  nature.  They  are  all  solids ; 
most  of  them  are  insoluble  in  water,  while  some  are 
insoluble  even  in  acids.  As  prepared  in  the  laboratory, 
some  of  these  salts,  such  as  copper  sulfide  (CuS)  and 
silver  sulfide  (Ag2S),  are  black ;  others,  as  cadmium  sul- 
fide (CdS)  and  arsenic  sulfide  (As2S3),  are  yellow;  while 
zinc  sulfide  (ZnS)  is  white. 

The  soluble  sulfides,  Na2S,  KaS,  (NH4)2S,  are  most  readily 
prepared  by  treating  the  appropriate  base  with  hydrosul- 
furic acid ;  the  insoluble  sulfides  may  be  prepared  by  heat- 
ing the  metals  with  sulfur,  although  the  general  and  more 
convenient  method  for  their  preparation  consists  in  passing 
hydrogen  sulfide  into  the  aqueous  solutions  of  appropriate 
salts  of  the  metals.  Thus,  copper  sulfide  may  be  prepared 
by  dissolving  copper  sulfate  (CuSO4)  in  water  and  passing 
hydrogen  sulfide  into  the  solution: 

H2S  +  CuS04  — >-  CuS  +  HSS04 


236    AN  ELEMENTABY  STUDY  OF  CHEMISTRY 

The  copper  sulfide,  being  insoluble,  precipitates  as  fast  as 
formed,  and  may  be  removed  from  the  liquid  by  nitration. 

The  preparation  of  these  sulfides  as  carried  out  in  the  labo- 
ratory may  be  illustrated  in  the  following  way  :  Hydrogen  sul- 
fide is  generated  in  a  Kipp  apparatus  A  (Fig.  96)  and  is  passed 
successively  into  bottles  B,  C,  D,  and  E,  containing,  respectively, 
the  aqueous  solutions  of  silver  nitrate,  cadmium  sulfate,  zinc 
acetate,  and  sodium  hydroxide.  As  the  gas  bubbles  through 


FIG.  96.   The  preparation  of  insoluble  sulfides  by  precipitation  with 
hydrogen  sulfide 

the  solutions  there  is  formed  black  silver  sulfide  (Ag2S)  in  £, 
yellow  cadmium  sulfide  (CdS)  in  C,  white  zinc  sulfide  (ZnS)  in  D. 
No  precipitate  is  produced  in  E,  for  although  sodium  sulfide  is 
formed,  it  is  soluble  in  water  and  therefore  does  not  separate.  • 

Oxides  of  sulfur.  Sulfur  forms  five  different  oxides. 
The  two  most  important  are  sulfur  dioxide  (SO2)  and 
sulfur  trioxide  (SO8).  Both  are  acid  anhydrides. 

Sulfur  dioxide  (sulfurous  anhydride)  (S02).  This  is  the 
well-known  gas  resulting  from  the  combustion  of  sulfur. 
It  occurs  in  nature  in  the  gas  issuing  from  volcanoes  and 
in  solution  in  the  waters  of  some  springs. 


SULFUR;  SELENIUM;   TELLURIUM  237 

Preparation.    Sulfur  dioxide  is  prepared  by  three  general 
methods  : 

1.  By  the  combustion  of  sulfur  or  a  metallic  sulfide.    In 

either  case  the  sulfur  is  converted  into  sulfur  dioxide  : 


S  +  0 


2  ZnS  -f  3  O2  -  >-  2  ZnO  +  2  SO2 

The  enormous  quantities  of  sulfur  dioxide  used  in  the 
manufacture  of  sulfuric  acid  are  prepared  by  this  general 
method. 

2.  By  the  reduction  of  sulfuric  acid.    When  concentrated 
sulfuric  acid  is  heated  with  certain  metals,  such  as  copper, 
a  part  of  the  acid  is  reduced  to  sulfurous  acid.    The  latter 
compound  then  decomposes  into  sulfur  dioxide  and  water, 
the  complete  equation  being  as  follows: 

Cu  +  2  H2SO4  —  *-  CuSO4  +  SO2  +  2  H2O 

3.  By  the  action  of  acids  upon  a  sulfite.    Sulfites  are  salts 
of  sulfurous  acid  (H2SOa).   When  an  acid,  such  as  hydro- 
chloric acid,  is  added  to  a  sulfite,  sulfurous  acid  is  formed, 
which   decomposes   into  water  and  sulfur   dioxide.     The 
reactions  are  expressed  in  the  following  equations: 

3  (1) 

(2) 

Explanation  of  the  reaction.  In  the  action  of  hydrochloric 
acid  upon  sodium  sulfite,  as  expressed  in  these  equations, 
we  have  two  reversible  reactions  depending  upon  each  other. 
It  might  be  expected  that  the  reaction  expressed  in  equa- 
tion (1)  would  result  in  an  equilibrium,  since  none  of  the 
substances  represented  in  the  equation  are  insoluble  or  vola- 
tile in  the  presence  of  water.  The  sulfurous  acid,  however, 
decomposes  as  fast  as  it  forms,  according  to  equation  (2), 
the  resulting  sulfur  dioxide  escaping  in  the  form  of  a  gas. 


238    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


The  reaction  continues,  therefore,  until  practically  all  the 
sodium  sulfite  has  been  decomposed.  Since  sulfur  dioxide  is 
quite  soluble  in  water,  it  is  evident  that  the  reaction  should 
be  carried  out  in  the  presence  of  as  little  water  as  possible  ; 
otherwise  a  proportionately  larger  quantity  of  sulfur  dioxide 
will  remain  in  solution,  and  the  reaction  will 
not  reach  the  same  degree  of  completion. 

Properties.  Sulfur  dioxide  is  a  colorless  gas 
which  at  ordinary  temperatures  is  2.2  times 
as  heavy  as  air.  It  has  a  peculiar,  irritating 
odor.  The  gas  is  very  soluble  in  water, 
1  volume  of  water  dissolving  approximately 
80  volumes  of  the  gas  under  standard  con- 
ditions. It  is  easily  condensed  to  a  colorless 
liquid  (boiling  point,  —  8°)  and  can  be  pur- 
chased in  this  condition,  stored  in  strong 
bottles  or  in  tin  cylinders  (Fig.  97). 

Chemical    conduct.     Sulfur    dioxide    has    a 
marked  tendency  to  combine  with  other  sub- 
stances and  is  therefore  an  active  substance 
chemically.     It    has    a    marked    affinity    for 
oxygen  and  is  therefore  a  reducing  agent.   Under  some  con- 
ditions it  can  also  act  as  an  oxidizing  agent.    Thus,  it  reacts 
with  hydrogen  sulfide  to  form  water  and  sulfur,  as  follows  : 


Since  both  hydrogen  sulfide  and  sulfur  dioxide  are  present 
in  the  gases  issuing  from  volcanoes,  it  is  probable  that  the 
large  deposits  of  sulfur  occurring  in  volcanic  regions  have 
resulted  from  the  interaction  of  these  two  gases,  according 
to  the  above  equation.  A  characteristic  property  of  sulfur 
dioxide  is  its  conduct  towards  water,  with  which  it  combines 
to  form  sulfurous  acid. 


SULFUR;  SELENIUM;  TELLURIUM          239 

Sulfurous  acid  (H2S03).  When  sulfur  dioxide  is  passed 
into  water  some  of  the  gas  combines  with  water  to  form 
sulfurous  acid  (H2SO3),  while  the  remainder  is  held  in  a 
state  of  solution.  The  sulfurous  acid  formed  is  in  equi- 
librium, on  the  one  hand,  with  water  and  dissolved  sulfur 
dioxide  and,  on  the  other  hand,  with  the  ions  H+  and 
HSO3~,  resulting  from  the  ionization  of  a  portion  of  the  acid: 

H2O  +  SO2  +=±  H2SO3  q=fc  H+  +  HSO,- 

When  heated  this  liquid  acts  as  if  it  were  simply  a 
solution  of  sulfur  dioxide  in  water,  all  the  sulfur  being 
evolved  as  sulfur  dioxide.  Toward  a  base,  on  the  other 
hand,  it  acts  simply  as  a  solution  of  sulfurous  acid  (com- 
pare with  aqua  ammonia,  p.  204). 

Because  of  its  unstable  character  sulfurous  acid  can 
be  obtained  only  in  the  form  of  a  dilute  solution.  This 
solution  has  the  following  properties: 

1.  Acid  properties.    The  solution  has  all  the  properties 
typical  of  a  very  weak  acid.    When  neutralized  by  bases 
sulfurous  acid  yields  a  series  of  salts  called  sulfites,  most 
of  which  are  insoluble  in  water. 

2.  Reducing  properties.    Solutions  of  sulfurous  acid  act 
as  good  reducing  agents.    This  is  due  to  the  fact  that  sul- 
furous acid  has  the  power  of  taking  up  oxygen  from  the 
air  or  from  substances  rich  in  oxygen,  and  is  changed  by 
this  reaction  into  sulfuric  acid: 

2H2S03  +  02 — ^2H2S04 

3.  Bleaching  properties.    Sulfurous  acid  has  strong  bleach- 
ing properties.    It  is  on  this  account  used  to  bleach  paper, 
straw  goods,  and  even  such  foods  as  canned  corn  and  dried 
fruits.    As  a  rule  the  bleaching  is  not  permanent.    It  is 
not  so  efficient  a  bleaching  agent  as  chlorine,    and  for 


240    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


this  reason  is  used  in  bleaching  only  such  materials  as  are 
injured  by  the  action  of  chlorine. 

The  bleaching  properties  of  sulfurous  acid  may  be  shown 
by  bringing  a  small  dish  of  burning  sulfur  under  a  bell  jar 
(Fig.  98)  in  which  has  been  placed  a  highly  colored  flower 
moistened  with  water.  Straw  hats  may  be  cleaned  and  bright- 
ened in  a  similar  way. 

4.  Antiseptic  properties.  Sulfurous  acid  has  marked  anti- 
septic properties,  and  on  this  account  has  the  power  of 
arresting  fermentation.  It  is  therefore 
used  in  certain  foods  containing  sugars, 
such  as  sweet  cider,  canned  corn,  and 
dried  fruits.  Whether  or  not  its  use  in 
foods  should  be  permitted  is  a  much 
debated  question. 

Salts  of  sulfurous  acid ;  sulfites.  Since 
it  contains  two  hydrogen  atoms,  sulfur- 
ous acid  forms  both  acid  and  normal 
salts.  Thus,  with  sodium  it  forms  the 
salts  NaHSO3  and  Na2SO3.  The  sulfites 
are  all  solid  substances  and,  like  sul- 
furous acid  itself,  combine  readily  with  oxygen,  forming  the 
corresponding  sulfates.  They  are  therefore  good  reducing 
agents.  Because  of  this  property,  unless  freshly  prepared, 
they  are  apt  to  contain  more  or  less  of  the  corresponding 
sulfates.  Calcium  acid  sulfite  is  used  in  the  manufacture 
of  paper  from  wood,  since  it  dissolves  the  objectionable 
constituent  (lignin)  of  the  wood,  leaving  the  pure  cellulose, 
which  is  the  material  desired  for  the  manufacture  of  paper. 
Sulfur  trioxide  (sulfuric  anhydride)  (S03).  When  sulfur 
is  burned  in  oxygen  minute  quantities  of  sulfur  trioxide 
are  formed  along  with  the  sulfur  dioxide.  Likewise,  when 


FIG.  98.    Bleaching 

a  flower  with  sulfur 

dioxide 


SULFUR;  SELENIUM;  TELLURIUM  241 

sulfur  dioxide  and  oxygen  are  heated  together,  combina- 
tion takes  place,  but  the  speed  of  the  reaction  is  so  slow 
that  only  traces  of  the  trioxide  result.  In  the  presence 
of  a  catalytic  agent,  however,  such  as  finely  divided  plati- 
num, the  speed  is  greatly  increased,  and  in  this  way  sul- 
fur trioxide  can  be  obtained  in  quantities.  The  reaction 
is  a  reversible  one,  as  is  shown  in  the  following  equation : 

2  SO2  4-  ( )2  +=>:  2  SO3  +  2  x  22,600  cal. 


FIG.  99.   The  preparation  of  solid  sulfur  trioxide 

The  largest  yield  of  sulfur  trioxide  is  obtained  when  the 
reaction  is  carried  out  at  approximately  400°,  at  which 
temperature  about  98  per  cent  of  the  sulfur  dioxide 
combines  with  oxygen. 

The  preparation  of  the  trioxide  by  the  last-named  method 
can  be  carried  out  in  the  laboratory  as  follows  :  The  platinum 
used  as  a  catalytic  agent  is  prepared  by  moistening  asbestos  fiber 
in,  a  solution  of  chloroplatinic  acid  and  igniting  it  in  a  flame, 
whereby  the  platinum  compound  is  reduced  to  metallic  plati- 
num. The  fiber  containing  the  finely  divided  platinum  is 
placed  in  a  tube  of  hard  glass  .1  (Fig.  99),  which  is  then 
heated  to  about  400°,  while  equal  volumes  of  sulfur  dioxide 
and  oxygen,  previously  dried  by  bubbling  them  through  sulfuric 
acid  (contained  in  bottles  B  and  C),  are  passed  into  the  tube. 


242    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

As  this  mixture  comes  in  contact  with  the  catalytic  agent  com- 
bination takes  place,  and  the  resulting  sulfur  trioxide  escapes 
from  the  jet  at  the  end  of  the  tube  and  may  be  condensed  by 
surrounding  the  receiving  tube  D  with  a  freezing  mixture. 

Properties  of  sulfur  trioxide.  Sulfur  trioxide  is  a  color- 
less liquid  which  solidifies  at  about  15°  and  boils  at  46°. 
A  trace  of  moisture  causes  it  to  solidify  into  a  mass  of 
silky  white  crystals  somewhat  resembling  asbestos  fiber 
in  appearance.  These  crystals  have  the  formula  (SO3)2, 
or  S2O6.  In  contact  with  the  air  sulfur  trioxide  fumes 
strongly,  and  when  thrown  upon  water  it  dissolves  with 
a  hissing  sound  and  the  liberation  of  a  great  deal  of  heat. 
The  product  of  this  reaction  is  sulfuric  acid,  so  that 
sulfur  trioxide  is  the  anhydride  of  that  acid: 

SO3  +  H2O >-H2SO4 

Sulfuric  acid  (H2S04).  Sulfuric  acid  has  long  been 
known,  and  was  one  of  the  most  important  reagents 
employed  by  the  alchemists.  It  is  by  far  the  most  largely 
used  of  all  the  acids.  Not  only  is  it  one  of  the  most 
common  reagents  in  the  laboratory,  but  enormous  quan- 
tities of  it  are  consumed  in  the  industries,  especially  in 
the  manufacture  of  fertilizers,  the  refining  of  petroleum, 
and  in  cleaning  scale  from  iron  and  steel. 

Manufacture  of  sulfuric  acid.  Two  general  methods  for 
the  manufacture  of  sulfuric  acid  are  in  use  at  the  present 
time.  These  are  known  as  the  contact  process  and  the 
lead-chamber  process. 

1.  Contact  process.  The  reactions  taking  place  in  this 
process  are  represented  by  the  following  equations : 

S  +  02— ^S02  (1) 

2SO2  +  O2 — >-2SO8  (2) 

S03  +  H20— >H2S04  (3) 


SULFUR;  SELENIUM;  TELLURIUM          243 

Sulfur  dioxide  is  prepared  according  to  equation  (1),  by 
burning  sulfur  or  some  sulfide,  such  as  iron  pyrite  (FeSg), 
in  air.  The  resulting  sulfur  dioxide,  together  with  suffi- 
cient air  to  furnish  the  necessary  oxygen,  is  conducted 
through  iron  tubes  filled  with  some  porous  material  (as- 
bestos or  sodium  sulfate),  through  which  a  suitable  cata- 
lytic agent,  such  as  platinum  or  iron  oxide,  is  interspersed, 
the  material  being  kept  at  about  400°.  Under  these  con- 
ditions sulfur  trioxide  is  formed  according  to  equation  (2). 
The  resulting  sulfur  trioxide  is  then  brought  into  contact 
with  water,  with  which  it  unites  to  form  sulfuric  acid 
according  to  equation  (3). 

The  only  part  of  the  process  which  is  difficult  to  carry  out 
on  a  commercial  scale  is  the  formation  of  the  sulfur  trioxide. 
It  has  long  been  known  that  sulfur  dioxide  and  oxygen  com- 
bine when  passed  over  finely  divided  platinum,  but  the  cost  of 
platinum,  together  with  the  poor  yield  of  sulfur  trioxide  ob- 
tained, made  the  process  an  impracticable  one.  A  study  of  the 
conditions  under  which  the  reaction  takes  place  resulted  in  im- 
provements in  the  process,  until  finally,  in  1901,  the  German 
chemist  Knietsch  succeeded  in  overcoming  the  difficulties  to 
such  an  extent  as  to  make  the  process  a  commercial  success  for 
the  manufacture  of  the  pure,  concentrated  acid.  While  plati- 
num is  the  most  effective  catalytic  agent  for  the  process,  it  is 
very  expensive,  its  commercial  value  at  the  present  time  being 
greater  than  that  of  gold.  This  has  led  to  the  use  of  other  cata- 
lytic agents,  among  which  iron  oxide  appears  to  be  the  best. 

It  is  an  interesting  fact  that  the  sulfur  .trioxide  produced  by 
this  method  will  not  combine  with  pure  water.  It  is  conducted 
into  concentrated  sulfuric  acid  and  combines  readily  with  the 
water  present  in  this  reagent. 

2.  Chamber  process.  The  older  method  of  manufacture, 
exclusively  employed  until  recent  years  and  still  the  most  im- 
portant process,  is  much  more  complicated.  The  conversion 


244    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

of  water,  sulfur  dioxide,  and  oxygen  into  sulfuric  acid  is 
accomplished  by  the  catalytic  action  of  oxides  of  nitrogen. 
Since  these  oxides  are  gases,  it  is  difficult  to  prevent  their 
escape,  and  very  elaborate  precautions  have  to  be  taken 
to  reduce  the  loss  as  much  as  possible.  The  reactions  are 
brought  about  in  large,  lead-lined  chambers,  into  which 
oxides  of  nitrogen,  sulfur  dioxide,  steam,  and  air  are  intro- 
duced in  suitable  proportions.  These  react  to  form  sulfuric 
acid,  which  collects  on  the  floor  of  the  chambers  and  is 
drawn  off  from  time  to  time. 

Reactions  of  the  chamber  process.  The  reactions  involved  are 
quite  complex,  and  are  not  at  all  thoroughly  understood.  It  is 
believed,  however,  that  the  two  following  general  reactions 
take  place :  (1)  The  substances  introduced  into  the  chambers 
first  react  to  form  a  derivative  of  sulfuric  acid  known  as  nitrosyl 
sulfuric  acid.  The  relation  of  these  two  compounds  to  each 
other  may  be  seen  from  their  structural  formulas : 

HO->«^°  H-O^Q^O 

HCT  b^O  NO-CK   ^O 

sulfuric  acid  nitrosyl  sulfuric  acid 

The  reaction  may  be  represented  as  follows  : 

TTO 

2  SO2  +  NO  +  NO2  +  H2O  +  O2 *•  2  ^Q  _  ^  >  SO2        (1) 

This  acid  can  be  obtained  in  the  form  of  white  crystals  known 
as  chamber  crystals. 

(2)  In  the  commercial  manufacture  of  sulfuric  acid,  however, 
such  a  separation  does  not  occur,  because  sufficient  water  is 
always  present  to  change  the  nitrosyl  acid,  as  fast  as  formed, 
into  sulfuric  acid : 

2  NQ!?o  >  S°2  +  H2° *  2  H2S°4  +  NO  +  JSTO2  (2) 

It  will  be  noted  that  in  equation  (2)  the  same  quantities  of  the 
oxides  of  nitrogen  are  formed  as  are  required  for  equation  (1). 
Theoretically,  therefore,  a  small  amount  of  these  oxides  should 


SULFUR;  SELENIUM;  TELLURIUM          245 

suffice  to  prepare  an  unlimited  amount  of  sulfuric  acid ;  practi- 
cally, some  of  the  oxides  are  lost,  and  this  loss  must  be  replaced. 
The  sulfuric  acid  plant.  The  simpler  parts  of  a  plant  used  in 
the  manufacture  of  sulfuric  acid  are  illustrated  in  Fig.  100. 
Sulfur  or  some  sulfide,  as  FeS2,  is  burned  in  the  furnace  A. 
The  resulting  sulfur  dioxide,  together  with  the  necessary 
amount  of  air,  passes  into  the  structure  C,  known  as  the  Glover 
tower.  In  it  the  oxides  of  nitrogen  are  generated,  as  will  be 
explained  later,  and  these,  together  with  the  sulfur  dioxide 
and  air,  pass  into  the  chambers  D,  D.  Water  or  steam  is  also 
introduced  into  these  chambers  at  suitable  points.  Here  the 


FIG.  100.    Diagram  of  a  plant  for  the  manufacture  of  sulfuric  acid 

reactions  take  place  which  result  in  the  formation  of  the  sul- 
furic acid.  The  nitrogen  remaining  after  the  withdrawal  of 
the  oxygen  from  the  air  which  entered  the  chamber  escapes 
through  the  structure  E,  known  as  the  Gay-Lussac  tower.  In 
order  to  prevent  the  escape  of  the  nitrogen  dioxide  regenerated 
in  the  reaction,  this  tower  is  filled  with  pieces  of  coke  over 
which  trickles  concentrated  sulfuric  acid  admitted  in  the  form 
of  a  spray  (F~)  at  the  top.  The  concentrated  acid  absorbs  the 
nitrogen  dioxide  but  not  the  nitric  oxide,  so  that  the  latter 
escapes  along  with  the  nitrogen.  The  acid  which  is  sprayed 
into  the  top  of  the  tower  collects  in  the  bottom  and  is  run  off 
into  the  vessel  G,  from  which  it  is  forced  into  the  tank  at  the 
top  of  the  Glover  tower  C.  Here  it  is  mixed  with  some  dilute 


246    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

sulfuric  acid,  and  the  mixture  sprayed  into  the  top  of  the 
tower,  which  is  partly  filled  with  some  acid-resisting  rock.  As 
the  acid  passes  down  through  this  material  it  meets  with  the 
hot  gases  entering  from  the  furnace,  whereby  the  nitrogen 
dioxide  is  liberated  from  the  acid,  passes  over  into  the  cham- 
bers D,  D,  and  again  enters  into  the  reaction.  During  the  process 
just  described  the  dilute  acid  becomes  sufficiently  concentrated 
to  serve  again  as  an  absorbent  of  nitrogen  dioxide.  The  neces- 
sary quantity  of  it  is  therefore  run  into  the  vessel  //  from  the 
bottom  of  the  tower,  and  then  forced  into  the  tank  at  the  top 
of  E.  In  order  to  replace  the  oxides  of  nitrogen  lost  in  the 
process,  the  necessary  quantity  is  added  by  the  action  of  sul- 
furic acid  upon  sodium  nitrate  in  vessel  B.  The  sulfuric  acid 
so  formed,  together  with  the  excess  of  condensed  steam,  collects 
upon  the  floor  of  the  chambers  in  the  form  of  a  liquid  contain- 
ing from  62  to  70  per  cent  of  hydrogen  sulfate.  The  product  is 
called  chamber  acid  and  is  quite  impure ;  but  for  many  pur- 
poses, such  as  the  manufacture  of  fertilizers,  it  needs  no 
further  treatment.  It  can  be  concentrated  by  evaporation  in 
vessels  variously  made  of  iron,  platinum,  or  silica. 

Relative  advantages  of  the  contact  process  and  the  lead-chamber 
process.  It  will  be  noted  that  in  the  contact  process  it  is  just 
as  easy  to  prepare  the  pure  concentrated  acid  as  the  dilute  acid. 
In  the  chamber  process,  however,  the  dilute  acid  is  obtained 
first  and  can  be  prepared  at  a  very  low  cost.  The  concentra- 
tion and  purification  of  the  dilute  acid  is,  however,  an  expen- 
sive operation.  For  these  reasons  the  contact  process  can 
compete  with  the  lead-chamber  process  only  in  the  manu- 
facture of  the  pure  concentrated  acid. 

Properties.  Pure  anhydrous  sulfuric  acid,  more  properly 
named  hydrogen  sulfate,  is  a  colorless,  oily  liquid  nearly 
twice  as  heavy  as  water.  The  ordinary  concentrated  acid 
contains  about  2  per  cent  of  water,  has  a  density  of  1.84, 
and  boils  at  338°.  It  is  sometimes  called  oil  of  vitriol, 
since  it  was  formerly  made  by  distilling  a  mixture  of 
substances,  one  of  which  was  called  green  vitriol. 


SULFUR;  SELENIUM;  TELLURIUM  247 

Chemical  conduct.  Sulfuric  acid  possesses  chemical  prop- 
erties which  make  it  one  of  the  most  important  of  chem- 
ical substances. 

1.  Acid  properties.     In  concentrated  aqueous  solutions, 
hydrogen  sulfate  forms  the  ions  H+  and  HSO4~,  the  latter, 
on  further  dilution  of  the  solution,  breaking  down  into  the 
ions  H+  and  SO4~~.    It  is  this  aqueous  solution  contain- 
ing hydrogen  ions  which  is  properly  termed  sulfuric  acid. 

2.  Action  as  an  oxidizing  agent.    Sulfuric  acid  contains  a 
large  percentage  of  oxygen  and  is,  like  nitric  acid,  a  very 
good  oxidizing  agent.  When  the  concentrated  acid  is  heated 
with  sulfur  or  carbon  or  various  other  substances,  oxidation 
takes  place,  the  sulfuric  acid  decomposing  according  to  the 
equation  H2SO4  — >  H2SO3  +  [O] 

3.  Action  on  metals.     A  dilute  solution  of  sulfuric  acid 
acts  upon  the  metals  that  precede  hydrogen  in  the  electro- 
chemical series  (p.  191),  forming  a  sulfate  of  the  metal 
and  hydrogen.     Such  a  solution  has  no  action  upon  the 
metals  that  follow  hydrogen  in  the  series. 

On  the  other  hand,  the  concentrated  acid  acts  upon  a 
number  of  the  metals  without  respect  to  their  position  in 
the  electrochemical  series ;  but  in  all  these  cases  the  first 
action  is  one  of  oxidation.  With  copper  the  reaction  is 
represented  by  the  equation 

Cu  +  H2SO4 >•  CuO  +  H2O  +  SO2 

The  copper  oxide  then  dissolves  in  an  additional  quantity 
of  sulfuric  acid  to  form  copper  sulfate : 

CuO  +  H2SO4  — >-  CuSO4  +  H3O 
These  two  equations  can  be  combined  into  the  form 
'  Cu  +  2  H2SO4  — h  CuSO4  +  2  HZO  +  SOa 


248    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

4.  Action  on  salts.     We  have  repeatedly  seen  that  an 
acid  of  high  boiling  point,  heated  with  the  salt  of  some 
acid    of    lower    boiling    point,    will    drive    out   the    low- 
boiling  acid  (p.  225).    The  boiling  point  of  sulfuric  acid 
(338°)   is  higher   than  that   of  almost  any  common  add; 
hence  it  is  largely  used  in  the  preparation  of  other  acids. 

5.  Action  on   water.     Concentrated  sulfuric  acid  has  a 
very  great  affinity  for  water,  and  when  mixed  with  water 
much  heat  is  evolved.    It  is  therefore  an  effective  drying, 
or  dehydrating,   agent.     Gases   which   have   no    chemical 
action  upon  sulfuric  acid  can  be  freed  from  water  vapor 
by  bubbling  them  through  the  concentrated  acid. 

Not  only  can  sulfuric  acid  absorb  water,  but  it  will  often 
withdraw  the  elements  hydrogen  and  oxygen  from  a  compound 
containing  them,  decomposing  the  compound  and  combining 
with  the  water  so  formed.  For  this  reason  most  organic  sub- 
stances, such  as  sugar,  wood,  cotton  and  "woolen  fiber,  and  even 
animal  tissues,  all  of  which  contain  much  oxygen  and  hydro- 
gen in  addition  to  carbon,  are  charred  by  the  action  of  the 
concentrated  acid.  The  process  in  general  consists  in  the  with- 
drawal of  the  oxygen  and  hydrogen  present  in  the  compound, 
thus  leaving  the  black  carbon  as  a  residue. 

Salts  of  sulfuric  acid ;  sulfates.  The  sulfates  constitute 
a  very  important  class  of  compounds,  and  many  of  them 
have  extensive  commercial  uses.  The  normal  salts  are 
all  solids  and,  with  the  exception  of  those  of  barium, 
strontium,  and  lead,  are  soluble  in  water.  Several  others, 
notably  calcium  sulfate  and  silver  sulfate,  are  only  slightly 
soluble. 

Other  oxygen  acids  of  sulfur.  In  addition  to  sulfurous  and 
sulfuric  acids  a  number  of  other  oxygen  acids  of  sulfur  are 
known,  either  in  the  free  state  or  in  the  form  of  their  salts. 
The  following  are  the  most  important : 


SULFUR;   SELENIUM;  TELLURIUM  249 

1.  Pyrosulfuric  acid  (fT2S2O7).  This  is  a  solid  crystalline  com- 
pound prepared  by  the  union  of  sulfuric  acid  and  sulfur 
trioxide  :  +  ^  -  ^  j^g^ 


The  acid  and  its  salts  are  strong  oxidizing  agents.  The 
fuming  sulfuric  acid  of  commerce  consists  of  a  mixture  of 
sulfuric  and  pyrosulfuric  acids. 

2.  Persulfuric  acid  (jET2S208).  This  acid  is  unstable  and  exists 
only  in  dilute  solution.  Its  salts,  however,  are  stable.  They  are 
prepared  by  the  electrolysis  of  concentrated  solutions  of  the 
corresponding  acid  sulfates.  Thus,  KHS04  by  electrolysis  yields 
K2S2Og.  The  salts  of  persulfuric  acid  are  very  strong  oxidiz- 
ing agents.  Thus,  ammonium  persulfate,  (NH4)2S2Og,  is  often 
used  as  an  oxidizing  agent,  especially  in  connection  with  certain 
photographic  processes. 

Monobasic  and  dibasic  acids.  Such  acids  as  hydrochloric 
and  nitric  acids,  whose  molecules  have  only  one  replaceable 
hydrogen  atom,  or,  in  other  words,  yield  one  hydrogen  ion 
in  solution,  are  called  monobasic  acids.  Acids  whose  mole- 
cules yield  two  hydrogen  ions  in  solution  are  called  dibasic 
acids.  Similarly,  we  may  have  tribasic  and  tetrabasic  acids. 
The  three  acids  of  sulfur  are  dibasic  acids.  It  is  therefore 
possible  for  each  of  them  to  form  both  normal  and  acid  salts. 

Preparation  of  acid  salts.  The  acid  salts  can  be  made 
in  two  ways:  the  acid  may  be  treated  with  only  half 
enough  base  to  neutralize  it,  — 

NaOH  +  H2SO4  —  >-NaHSO4  +  H2O 
or  a  normal  salt  may  be  treated  with  the  free  acid,  — 

Na2SO4  +  H2S04  —  +  2  NaHSO4 

Carbon  disulfide  (CS2).  When  sulfur  vapor  is  passed 
over  highly  heated  carbon,  the  two  elements  combine, 
forming  carbon  disulfide  : 

C  +  2  S  —  +  CS2  -  19,600  cal. 


250    AK  ELEMENTARY  STUDY  OF  CHEMISTRY 


Since  heat  is  absorbed  in  this  reaction,  it  must  be  supplied 
from  external  sources,  and  the  reaction  will  proceed  only 
at  a  rather  high  temperature. 

Carbon  disulfide  is  a  heavy,  colorless,  highly  refractive 
liquid  which  boils  at  46°.  When  pure  it  has  a  pleasant 
odor,  but  it  gradually  undergoes  slight  decomposition  and 
acquires  a  most  disagreeable  odor. 
Its  vapor  is  very  inflammable,  burn- 
ing in  the  air  to  form  carbon  dioxide 
and  sulfur  dioxide : 

CS2  +  3  O2 >-  CO2  +  2  SO2 

Carbon  disulfide  is  a  good  solvent 
for  many  substances,  such  as  gums, 
resins,  and  waxes,  which  are  not 
soluble  in  most  liquids,  and  it  is 
therefore  used  as  a  solvent  for  such 
substances.  It  is  also  used  as  an 
insecticide.  Its  vapor  is  poisonous 
as  well  as  highly  inflammable,  so 
that  one  must  exercise  great  care  in 
working  with  it. 


FIG.  101.   An  electric  fur- 
nace for  the  production  of 
carbon  disulfide 


Commercial  preparation  of  carbon  di- 
sulfide. Commercially,  carbon  disulfide 
is  made  by  the  direct  combination  of 

carbon  and  sulfur,  the  heat  necessary  for  this  union  being  de- 
rived from  an  electric  current.  The  main  part  of  a  large  fur- 
nace A  (Fig.  101)  is  filled  with  charcoal  introduced  through  the 
trap  C.  Sulfur  is  added  through  the  hoppers  D,  D.  An  electric 
current  is  passed  in  at  E,  E.  The  heat  generated  is  sufficient 
to  vaporize  the  sulfur,  which  then  unites  with  the  hot  carbon 
to  form  carbon  disulfide.  The  vapors  escape  at  H  and  are  con- 
densed. Some  of  the  furnaces  are  40  ft.  in  height  and  yield 
as  much  as  25,000  Ib.  of  the  disulfide  in  twenty-four  hours. 


SULFUR;  SELENIUM;  TELLURIUM          251 

Selenium  and  tellurium.  These  two  rather  uncommon  elements 
are  closely  related  to  sulfur  in  their  chemical  conduct.  They 
are  usually  found  associated  with  sulfur  and  sulfides,  either 
as  the  free  elements  or,  more  commonly,  in  combination  with 
metals.  With  hydrogen  they  form  compounds  of  the  formulas 
H2Se  and  H2Te ;  these  bodies  are  gases  with  properties  very 
similar  to  those  of  H2S.  They  also  form  oxides  and  oxygen 
acids  which  resemble  the  corresponding  sulfur  compounds. 
The  elements  even  have  forms  corresponding  very  closely  to 
those  of  sulfur. 

The  hydrates.  Attention  has  been  called  to  the  fact 
that  many  compounds  combine  directly  with  water,  form- 
ing new  compounds  known  as  hydrates.  Many  salts  possess 
this  property  in  a  marked  degree,  forming  hydrated  salts. 
Thus,  when  copper  sulfate  (CuSO4)  is  dissolved  in  water 
and  the  solution  is  set  aside,  a  crystalline  hydrate  is  de- 
posited which  is  formed  by  the  union  of  1  molecule  of 
copper  sulfate  with  5  molecules  of  water,  and  its  formula 
is  written  CuSO4  •  5  H2O.  Commercially,  this  hydrate  is 
called  blue  vitriol.  Similarly,  magnesium  sulfate  (MgSO4) 
forms  the  crystalline  hydrate  MgSO4  •  7  H2O.  If  we  wish 
to  distinguish  between  the  salt  and  its  hydrate,  the  salt 
is  referred  to  as  the  anhydrous  compound.  Thus,  the 
salt  CuSO4  is  termed  anhydrous  copper  sulfate.  The 
hydrates  are  true  chemical  compounds,  any  given  hydrate 
being  formed  by  the  union  of  definite  weights  of  the 
anhydrous  compound  and  of  water.  Many  anhydrous 
compounds,  however,  combine  with  several  different  per- 
centages of  water  to  form  different  hydrates.  The  hydrates 
as  a  rule  are  not  very  stable  and  tend  to  decompose  into 
the  constituents  from  which  they  are  formed,  especially 
when  heated.  On  the  other  hand,  many  anhydrous  salts, 
such  as  calcium  chloride,  absorb  moisture  from  the  air 
and  form  hydrates. 


252    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Efflorescence.  All  hydrates  may  be  regarded  as  tending 
to  lose  water  at  ordinary  temperatures,  the  reaction  being 
reversible;  thus,  . 

Na2SO4 .  10  H2O  ^=±Na2SO4  +10  II2() 

If  crystals  of  a  hydrate  are  placed  in  a  closed  space,  the 
reaction  goes  on  until  the  pressure  of  the  water  vapor 
formed  reaches  a  definite  value,  and  then  equilibrium 
results.  If  this  pressure  is  greater  than  the  pressure  of 
water  vapor  in  the  air,  the  crystal  will  lose  water  when 
exposed  to  the  air  and  crumble  to  powder.  Such  crystals 
are  said  to  be  efflorescent.  Hydrated  sodium  sulfate  is  a 
good  example  of  an  efflorescent  compound. 

EXERCISES 

1.  Is  the  equation  for  the  preparation  of    hydrogen  sulfide   a 
reversible  one?   As  ordinarily  carried  out,  does  the  reaction  com- 
plete itself  ? 

2.  Suppose  that  hydrogen  sulfide  were  a    liquid ;    would  it  be 
necessary  to  modify  the  method  of  preparation  ? 

3.  Does  perfectly  dry  hydrogen  sulfide  change  the  color  of  litmus 
paper  ?   State  reason  for  your  answer. 

4.  What   is  an  acid    anhydride?    Aside   from  those  of   sulfur, 
what  other  anhydrides  have  been  mentioned? 

5.  How  would  you  expect  dilute  sulfuric  acid  to  act  upon  iron  ? 
upon  silver  ?    (Refer  to  electrochemical  series.) 

6.  Can  you  suggest  a  reason  why  silver  spoons  become  tarnished 
when  in  contact  with  certain  kinds  of  food  ? 

7.  Mention  other  instances  of  catalysis  aside  from  those  given 
in  this  chapter. 

8.  In  the  commercial  preparation  of  carbon  disulfide  what  is 
the  function  of  the  electric  current? 

9.  Write  the  equation  representing  the  reaction  between  hydro- 
sulfuric  acid  and  sodium   hydroxide;   between  hydrosulfuric  acid 
and  ammonium  hydroxide. 


SULFUR;  SELENIUM;  TELLURIUM          253 

10.  Show  that  the  preparation  of  sulfur  dioxide  from  a  sulfite 
is  similar  in  principle  to  the  preparation  of  hydrogen  sulfide. 

11.  Calculate  the  weight  of  materials  necessary  for  the  prepara- 
tion of  sufficient  hydrogen  sulfide  to  saturate  20  1.  of  water  at  15° 
and  normal  pressure. 

12.  What  weight  of  sulfur  is  necessary  for  the  preparation  of 
2000  Ib.  of  sulfuric  acid  containing  5  per  cent  of  water  ? 

13.  Suppose  you  wish  to  prepare  100  kg.  of  blue  vitriol ;  calcu- 
late the  weights  of  materials  necessary  for  its  preparation. 

14.  50  g.  of  blue  vitriol  was  dissolved  in  water,  and  hydrogen  sul- 
fide passed  through  the  solution  until  the  copper  was  all  precipitated. 
Calculate  the  weight  of  the  precipitate. 

15.  Write  the  names  and  formulas  of  the  oxides  and  oxygen  acids 
of  selenium  and  tellurium. 

16.  Contrast  the  action  of  dilute  sulfuric  acid  and  of  concen- 
trated sulfuric  acid  upon  zinc. 

17.  How  many  calories  of  heat  are  absoi'bed  in  the  preparation 
of  10  kg.  of  carbon  disulfide  ? 

18.  Write  equations  for  the  preparation  of  sodium  acid  sulfite  by 
two  different  methods. 


CHAPTER  XXI 
THE  PERIODIC  LAW 

A  number  of  the  elements  have  now  been  studied  some- 
what closely.  Of  these,  oxygen,  hydrogen,  nitrogen,  and 
chlorine,  while  having  some  physical  properties  in  common 
with  each  other,  have  almost  no  points  of  similarity  in 
their  chemical  conduct.  On  the  other  hand,  oxygen  and 
sulfur,  while  quite  different  physically,  have  much  in 
common  in  their  chemical  properties. 

More  than  eighty  elements  are  now  known.  If  all  of 
these  should  have  properties  as  diverse  as  do  oxygen, 
hydrogen,  nitrogen,  and  chlorine,  the  study  of  chemistry 
would  plainly  be  very  difficult  and  complicated.  If, 
however,  the  elements  can  be  classified  in  groups  the 
members  of  which  have  very  similar  properties,  the  study 
will  be  very  much  simplified. 

Earlier  classification  of  the  elements.  Even  at  an  early 
period  efforts  were  made  to  discover  some  natural  principle 
in  accordance  with  which  the  elements  could  be  classified. 
Two  of  these  classifications  may  be  mentioned  here. 

1.  Classification  into  metals  and  nonmetals.  The  classifica- 
tion into  metals  and  nonmetals  most  naturally  suggested 
itself.  This  grouping  was  based  largely  on  physical  proper- 
ties, the  metals  being  heavy,  lustrous,  malleable,  ductile,  and 
good  conductors  of  heat  and  electricity.  Elements  possess- 
ing these  properties  are  usually  base-forming  in  character, 
and  the  ability  to  form  bases  came  to  be  regarded  as  a 
254 


THE  PERIODIC  LAW 


255 


characteristic  property  of  the  metals.  The  nonmetals  pos- 
sessed physical  properties  which  were  opposite  to  those  of 
the  metals,  and  were  acid-forming  in  character. 

Not  much  was  gained  by  this  classification,  and  it  was  very 
imperfect.  Some  metals,  as  potassium,  are  very  light ;  some 
nonmetals,  as  iodine,  have  a  high  luster :  some  elements  can 
form  either  an  acid  or  a  base. 

2.  Classification  into  triad 
families.  In  1825  Dobereiner 
observed  that  an  interesting 
relation  exists  between  the 
atomic  weights  of  chemically 
similar  elements.  To  illus- 
trate, lithium,  sodium,  and  po- 
tassium resemble  each  other 
very  closely,  and  the  atomic 
weight  of  sodium  is  almost 
exactly  an  arithmetical  mean 
between  those  of  the  other 
6.94  +  39.10 


two:    - 


=  23.02. 


FIG.  102.    Mendele'eff  (1834-1907) 


A  Russian  chemist  who  proposed  the 
periodic  classification  of  the  elements 


Iii  many  chemical  and  physi- 
cal properties  sodium  is  mid- 
way between  the  other  two. 

A  number  of  triad  families  were  found,  but  among 
eighty  elements,  whose  atomic  weights  range  all  the  way 
from  1  to  240,  such  agreements  might  be  mere  chance. 
Moreover,  many  elements  did  not  appear  to  belong  to 
such  families. 

Periodic  classification.  In  1869  the  Russian  chemist 
Mendele'eff  (Fig.  102)  devised  an  arrangement  of  the  ele- 
ments based  on  their  atomic  weights  which  has  proved  to  be 
of  great  service  in  the  comparative  study  of  the  elements. 


256    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

A  few  months  later  the  German,  Lothar  Meyer,  independ- 
ently suggested  the  same  ideas.  This  arrangement  brought 
to  light  a  great  generalization,  now  known  as  the  periodic 
law.  An  exact  statement  of  the  law  will  be  given  in  a  later 
paragraph. 

Arrangement  of  the  periodic  table.  The  general  arrange- 
ment suggested  by  Mendeleeff  and  extended  so  as  to  include 
elements  more  recently  discovered  is  as  follows :  Omitting 
hydrogen,  which  has  the  smallest  atomic  weight,  and  begin- 
ning with  helium,  which  has  an  atomic  weight  of  3.99,  the 
first  eight  elements  are  arranged  in  a  horizontal  row  in 
the  order  of  their  atomic  weights.  These  eight  elements 
all  differ  markedly  from  each  other,  but  the  one  having 
the  next  highest  atomic  weight,  neon,  is  very  similar  to 
helium.  It  is  placed  just  under  helium,  and  a  new  hori- 
zontal row  follows  as  shown  below.  The  element  following 
chlorine,  namely  argon,  resembles  helium  and  neon  and 
begins  a  third  row.  These  three  rows  are  as  follows: 

He  (3.99)  Li  (6.94)  Gl  (9.1)       B  (11)      C  (12.005)  N  (14.01)  O  (16)      F  (19) 
Ne  (20.2)  Na  (23)    Mg  (24.32)  Al  (27.1)  Si  (28.3)     P  (31.04)   S  (32.06)  Cl  (35.46) 
A  (39.88)  K  (39.1)  Ca  (40.07)  Sc  (44.1)  Ti  (48.1)    V  (51)       Cr  (52)     Mn  (54.93) 

If  now  we  consider  the  elements  that  fall  into  vertical 
columns  in  these  three  rows,  a  remarkable  fact  is  brought 
to  light.  Not  only  is  there  a  strong  similarity  between 
helium,  neon,  and  argon,  which  form  the  first  vertical 
column,  but  the  elements  in  the  other  columns  exhibit 
much  of  the  same  kind  of  similarity  among  themselves, 
and  evidently  form  natural  groups.  Thus  lithium,  sodium, 
and  potassium  resemble  each  other  very  closely  and  form 
one  of  Dobereiner's  triads. 

Iron,  cobalt,  and  nickel,  following  manganese  (Mn), 
have  atomic  weights  near  together,  and  are  very  similar 
chemically.  They  do  not  strongly  resemble  any  of  the 


THE  PERIODIC  LAW  257 

elements  so  far  considered,  and  so  are  placed  in  a  group 
by  themselves.  The  first  three  horizontal  rows  of  the 
table  on  the  next  page  show  the  arrangement  of  these 
twenty-seven  elements.  A  new  horizontal  row  is  begun 
with  copper.  Following  the  fifth  and  seventh  rows  are 
groups  of  three  closely  related  elements,  so  the  completed 
arrangement  has  the  appearance  represented  in  the  table. 

The  relation  of  properties  of  elements  to  atomic  weights. 
There  is  evidently  a  close  relation  between  the  properties 
of  an  element  and  its  atomic  weight.  For  example,  con- 
sider the  elements  in  the  first  horizontal  row.  Helium 
is  an  inert  element.  Folio  whig  it,  lithium  is  a  metallic 
element,  has  a  valence  of  1,  and  possesses  a  strong  base- 
forming  character.  The  next  element,  glucinum,  has  a 
valence  of  2,  and  is  less  strongly  base-forming,  while 
boron  has  some  base-forming  and  some  acid-forming 
properties.  In  carbon,  all  base-forming  properties  have 
disappeared,  and  the  acid-forming  properties  are  more 
marked  than  in  boron.  These  become  still  more  empha- 
sized as  we  pass  through  nitrogen  and  oxygen,  until 
on  reaching  fluorine  we  have  one  of  the  strongest  acid- 
forming  elements.  The  properties  of  these  eight  elements 
vary  regularly  with  their  atomic  weights,  or,  in  mathe- 
matical language,  are  continuous  functions  of  them. 

The  periodic  law.  If  it  were  true  that  helium  had  the 
smallest  atomic  weight  of  any  of  the  elements  and  fluorine 
the  greatest,  so  that  in  passing  from  one  to  the  other  we 
included  all  the  elements,  we  could  say  that  the  properties 
of  elements  were  continuous  functions  of  their  atomic 
weights.  But  fluorine  is  an  element  of  relatively  small 
atomic  weight,  and  the  one  following  it,  neon,  breaks 
the  regular  order,  for  in  it  reappear  all  the  characteristic 
properties  of  helium.  Sodium,  following  neon,  bears  much 


258    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


811 

""ii"ii 


THE  PERIODIC  LAW  259 

the  same  relation  to  lithium  that  neon  does  to  helium,  and 
the  properties  of  the  elements  in  the  second  row  vary 
much  as  do  the  properties  of  the  elements  in  the  first 
row,  until  argon  is  reached,  when  another  repetition  be- 
gins. The  properties  of  the  elements  do  not  vary  con- 
tinuously, therefore,  with  atomic  weights,  but  at  regular 
intervals  there  is  a  repetition,  or  period.  This  general- 
ization is  known  as  the  periodic  law,  and  may  be  stated 
thus :  The  properties  of  elements  are  periodic  functions  of 
their  atomic  weights. 

Two  families  in  a  group.  The  elements  of  each  group 
(excepting  Group  0)  fall  naturally  into  two  families. 
The  elements  in  the  odd-numbered  horizontal  rows 
(called  periods')  form  one  family,  those  in  the  even- 
numbered  periods,  the  other.  In  the  table  these  are 
arranged  under  the  headings  A  and  B.  The  elements  in 
one  family  are  much  more  similar  to  each  other  than  they 
are  to  those  in  the  other  family  in  the  same  group.  Thus, 
magnesium,  zinc,  cadmium,  and  mercury  form  one  family 
of  very  similar  elements  in  Group  II,  while  glucinum, 
calcium,  strontium,  barium,  and  radium  form  the  other. 

Family  resemblances.  Let  us  inquire  more  closely  in. 
what  respects  the  elements  of  a  family  resemble  each  other. 

1.  Valence.  In  general  the  valence  of  the. elements  in 
a  family  is  the  same,  and  the  formulas  of  their  compounds 
are  therefore  similar.  If  we  know  that  the  formula  of 
sodium  chloride  is  NaCl,  it  is  pretty  certain  that  the 
formula  of  potassium  chloride  will  be  KC1  —  not  KC12  nor 
KClg.  The  general  formulas  R2O,  RO,  etc.,  placed  below 
the  columns,  indicate  the  formulas  of  the  oxides  of  the  ele- 
ments in  the  column,  provided  they  form  oxides.  In  like 
manner  the  formulas  RH,  RH2,  etc.  show  the  composition 
of  the  compounds  formed  with  hydrogen  or  with  chlorine. 


260    AN  ELEMENTARY  STUDY  OF  CHEMISTEY 

2.  Chemical  conduct.    The  chemical  characteristics  of  the 
members  of  a  family  are  quite  similar.    If  one  member  is 
a  metal,  the  others  usually  are ;  if  one  is  a  nonmetal,  so, 
too,  are  the  others.    The  families  in  Group  I  and  Group  II 
consist  of  metals,  while  the  elements  found  in  Group  VI 
and  Group  VII  form  acids.    There  is  in  addition  a  certain 
regularity  in  properties  of  the  elements  in  each  family. 
If  the  element  at  the  head  of  the  family  is  a  strong  acid- 
forming  element,  this  property  is  likely  to  diminish  gradu- 
ally, as  we  pass  to  the  members  of  the  family  with  higher 
atomic  weights.    Thus,  phosphorus  is  strongly  acid-forming, 
arsenic  less  so,  and  antimony  still  less  so,  while  bismuth 
has  almost  no  acid-forming  properties.    As  a  result  of  this 
regularity,   the    elements    of   high   atomic   weight   at   the 
bottom  of  the   family  columns  are  nearly  all  metals  in 
their  chemical  conduct.    We  shall  meet  with  many  illus- 
trations of  this  fact. 

3.  Physical  properties.     In  the  same  way,  the  physical 
properties  of  the  members  of  a  family  show  a  gradation 
as  we  pass  from  element  to  element  in  the  family.    Thus, 
the  densities  of  the  members  of  the  magnesium  family  are 

Mg  =  1.75,  Zn  =  7.00,  Cd  =  8.67,  Hg  =  13.6 
Their  melting  points  are 

Mg  =  651°,  Zn  =  419.4°,  Cd  =  320.9°,  Hg  =  -38.9° 

Value  of  the  periodic  law.  The  periodic  law  has  proved 
of  much  value  in  the  development  of  the  science  of 
chemistry. 

1.  It  simplifies  study.  It  is  at  once  evident  that  such 
regularities  very  much  simplify  the  study  of  chemistry. 
A  thorough  study  of  one  element  of  a  family  makes  the 
study  of  the  other  members  a  much  easier  task,  since  so 


THE  PERIODIC  LAW 


261 


many  of  the  properties  and  chemical  reactions  of  the 
elements  are  similar.  Thus,  having  studied  the  element 
sulfur  in  some  detail,  it  is  not  necessary  to  study  selenium 
and  tellurium  so  closely,  for  most  of  their  properties  can 
be  predicted  with  a  fair  degree  of  accuracy  from  the 
relation  which  they  sustain  to  sulfur. 

2.  It  predicts  new  elements.  When  the  periodic  law  was 
first  formulated  there  were  several  vacant  places  in 
the  table  which  evidently  belonged  to  elements  at  that 
time  unknown.  From  their  position  in  the  table,  Men- 
deleeff predicted  with  great  precision  the  properties  of  the 
elements  which  he  felt  sure  would  one  day  be  discovered  to 
fill  these  places.  Three  such  elements  —  scandium,  germa- 
nium, and  gallium  —  were  found  within  fifteen  years,  and 
their  properties  agreed  in  a  remarkable  way  with  the  pre- 
dictions of  Mendeleeff.  This  is  shown  in  the  following 
table,  in  which  the  properties  of  gallium  are  compared 
with  those  which  Mendeleeff  predicted: 


PROPERTIES  OF  GALLIUM 

PREDICTED 

FOUND 

Atomic  weight  
Melting  point     

about  69 
low 

69.9 
30° 

Specific  gravity.      
Formula  of  oxide  

5.9 
R0Oa 

5.95 
Ga0Oo 

Action  of  air     

no  action 

f  only  slight,  even 
\     at  red  heat 

According  to  the  table  there  are  still  some  undiscov- 
ered elements,  and  much  effort  has  been  made  to  find 
them.  In  the  family  with  manganese  there  should  be  an 
element  of  atomic  weight  approximately  100,  and  another 
of  about  187.  There  are  also  two  vacant  places  in  the 
last  period. 


262    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

3.  It  indicates  probable  errors.  The  physical  constants  of 
many  of  the  elements  did  not  at  first  agree  with  those 
demanded  by  the  periodic  law,  and  a  further  study  of 
many  such  cases  showed  that  errors  had  been  made. 
The  law  has  therefore  done  much  service  in  indicating 
probable  error. 

Imperfections  of  the  periodic  table.  There  still  remain 
a  good  many  features  which  must  be  regarded  as  im- 
perfections in  the  table.  Most  conspicuous  is  the  fact 
that  the  element  hydrogen  has  no  place  in  it.  In  some 
of  the  groups  elements  appear  in  one  of  the  families, 
while  all  of  their  properties  show  that  they  belong  in 
the  other.  Thus,  sodium  belongs  with  lithium  and  not 
with  copper;  fluorine  belongs  with  chlorine  and  not  with 
manganese.  There  are  three  instances  where  the  elements 
must  be  transposed  in  order  to  make  them  fit  into  their 
proper  group.  According  to  their  atomic  weights,  tellu- 
rium should  follow  iodine,  argon  should  follow  potassium, 
and  nickel  should  follow  cobalt.  Their  properties  show 
in  each  case  that  this  order  must  be  reversed.  The  table 
separates  some  elements  altogether  which  in  many  respects 
have  closely  agreeing  properties.  Iron,  chromium,  and  man- 
ganese are  in  different  groups,  although  they  are  similar 
in  many  respects. 

The  system  is  therefore  to  be  regarded  as  but  a  partial 
and  imperfect  expression  of  some  very  important  and 
fundamental  relation  between  the  substances  which  we 
know  as  elements,  the  exact  nature  of  this  relation  being 
as  yet  not  completely  clear  to  us.  The  knowledge  gained 
in  recent  years  from  the  study  of  radium  and  similar 
elements  (Chapter  XLII)  has  given  us  some  insight  into 
the  nature  of  this  relationship,  but  it  would  take  us  too 
far  to  pursue  the  matter  to  greater  detail. 


THE  PERIODIC  LAW  263 

EXERCISES 

1.  Suppose  that  an  element  were  discovered  that  filled  the  blank 
in  Group  0,  Period  4 ;  what  properties  would  it  probably  have  ? 

2.  Suppose  that  an  element  were  discovered  that  filled  the  blank 
in  Group  VI,  Period  8,  family  B ;  what  properties  would  it  have  ? 

3.  Sulfur  and  oxygen  both  belong  in  Group  VI,  although  in 
different  families;  in  what  respects  are  the  two  similar? 

4.  Consult  encyclopedia  for  some  of  the  notable  events  in  the 
life  of  Mendeleeff. 

5.  What  processes  of  nature  can  you  think  of  that  you  would 
call  periodic? 

6.  Can  you  think  of  any  mechanical  processes  that  illustrate  a 
periodic  function? 


CHAPTER  XXII 
THE  CHLORINE  FAMILY 


NAME 

ATOMIC 

WEIGHT 

MELTING 

POINT 

BOILING 

POINT 

COLOR  AND  STATE 

Fluorine  (F)     .     . 

19.00 

-223° 

-  187° 

pale-yellowish  gas 

Chlorine  (Cl)  •  .     . 

35.46 

-  101.5° 

-  33.6° 

greenish-yellow  gas 

Bromine  (Br)    .     . 

79.92 

-7.3° 

63° 

red  liquid 

Iodine  (I)     ... 

126.92 

113.5° 

184.4° 

purplish-black  solid 

NOTE.  Chlorine,  as  well  as  hydrogen  chloride  and  hydrochloric  acid, 
has  been  described  in  Chapter  XIV,  and  that  chapter  should  be  reviewed 
in  connection  with  the  general  study  of  the  chlorine  family. 

The  family.  The  four  elements  named  in  the  above 
table  form  a  strongly  marked  family  of  elements  and 
illustrate  very  clearly  the  way  in  which  the  members  of 
a  family  in  a  periodic  group  resemble  each  other,  as  well 
as  the  character  of  the  differences  which  we  may  expect 
to  find  between  the  individual  members. 

1.  Occurrence.    These  elements  do  not  occur  in  nature 
in  the  free  state.    The  compounds  of  the  last  three  ele- 
ments of  this  family  are  found  extensively  in  sea  water, 
and  on  this  account  the  name  halogens,  signifying  "  pro- 
ducers of  sea  salt,"  is  sometimes  applied  to  the  family. 

2.  Properties.    By  reference  to  the  table  it  will  be  seen 
that  the  melting  point  and  the  boiling  point  of  the  ele- 
ments of  this  family  increase  with  their  atomic  weights. 
A  somewhat  similar  gradation  is  noted  in  their  color  and 
state.    The  affinity  of  these  elements  for  hydrogen  and  the 

264 


THE  CHLORINE  FAMILY  265 

metals  is  in  the  inverse  order  of  their  atomic  weights, 
fluorine  having  the  strongest  affinity  and  iodine  the 
weakest.  Only  chlorine  and  iodine  form  oxides,  and  those 
of  the  former  element  are  very  unstable. 

3.  Compounds  with  hydrogen.  Hydrogen  combines  with 
each  of  the  elements  of  the  family  to  form  the  following 
important  hydrides : 

Hydrogen  fluoride  (H2F2)  :  a  colorless  liquid  boiling  at  19.4°. 
Hydrogen  chloride  (HC1)  :  a  colorless  gas  liquefying  at  —  83.1°. 
Hydrogen  bromide  (HBr)  :  a  colorless  gas  liquefying  at  —  73°. 
Hydrogen  iodide  (HI)  :  a  colorless  gas  liquefying  at  —  34.1°. 

In  the  complete  absence  of  water  these  hydrides  are 
rather  inactive  and  have  neither  acid  nor  basic  properties. 
They  dissolve  in  water,  however,  forming  solutions  that 
are  acid  in  character. 

FLUORINE 

Occurrence.  Fluorine  occurs  in  nature  most  abundantly 
in  the  mineral  fluorite  (CaF2),  in  cryolite  (Na3AlF6),  and 
in  the  complex  mineral  fluorapatite,  3  Ca3(PO4)2  •  GaF2. 
Traces  of  compounds  of  fluorine  are  also  found  in  many 
other  minerals,  in  sea  water,  in  bones,  and  in  the  enamel 
of  the  teeth. 

Preparation.  While  the  compounds  of  fluorine  have 
been  known  for  a  long  time,  all  attempts  to  isolate  the 
free  element  resulted  in  failure  until  1886,  when  the 
French  chemist  Moissan  (Fig.  103)  succeeded  in  obtain- 
ing it  in  a  pure  state.  He  prepared  it  by  the  electrolysis 
of  hydrogen  fluoride  in  which  had  been  dissolved  a  little 
potassium  hydrogen  fluoride  (KHF2)  to  render  the  liquid 
an  electrolyte.  The  solution  was  placed  in  a  U-shaped  tube 
(Fig.  104)  made  of  platinum  (or  copper),  which  was  fur- 
nished with  electrodes  and  exit  tubes  for  the  escape  of 


266    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

the  fluorine  and  hydrogen  which  are  liberated,  the  former 
at  the  anode  and  the  other  at  the  cathode.  Since  hydrogen 
fluoride  boils  at  19.4°,  the  process  must  be  carried  out  at  a 
low  temperature  to  prevent  the  liquid  from  vaporizing. 

Properties  and  conduct.  Fluorine  is  a  gas,  slightly 
yellow  in  color,  and  is  1.3  times  as  heavy  as  air.  It  can 
be  obtained  in  the  form  of  a  yellow  liquid  which  boils  at 


FIG.  103.   Tablet  erected  by  the  associates  and  friends  of  Moissan,  in 

his  laboratory  in  Paris,  in  1906,  on  the  twentieth  anniversary  of  the 

isolation  of  fluorine 

-  187°  and  solidifies  at  -  223°.  Chemically,  it  is  one  of 
the  most  active  of  all  elements.  Most  of  the  elements, 
when  brought  in  contact  with  fluorine,  combine  with  it 
with  so  much  energy  as  to  produce  light.  It  unites  with 
hydrogen  with  explosive  violence  and  readily  abstracts  it 
from  its  compounds.  For  example,  it  decomposes  water 
with  great  energy,  forming  hydrogen  fluoride  and  oxygen : 

2  F2  +  2  H2O  — *  2  H2F2  +  O2  +  2  x  6800  cal. 


THE  CHLORINE  FAMILY 


267 


It  liberates  all  the  other  members  of  the  chlorine  family 
from  their  compounds  with  hydrogen  and  the  metals. 
It  does  not  combine  with  oxygen,  however,  and  but 
superficially  attacks  gold,  platinum,  or  copper. 

Hydrogen  fluoride  (H2F2).  Hydrogen  fluoride  is  readily 
obtained  from  fluorite  by  the  action  of  concentrated  sulfuric 
acid.  The  equation  is  as  follows : 

CaF2  +  H2S04  — ^  CaS04  +  H,Ft 

The  formula  is  usually  written  H2F2, 
although  by  selecting  the  proper  tem- 
perature the  compound  may  be  ob- 
tained in  any  of  the  forms  indicated 
by  the  formulas  HF,  H2F?,  H8F3.  In 
its  properties  hydrogen  fluoride  resem- 
bles the  hydrides  of  the  other  elements 
of  this  family,  although  it  is  more 
easily  condensed  to  a  liquid.  It  boils 
at  19.4°  and  can  therefore  be  liquefied 
at  ordinary  pressures.  It  is  soluble 
in  all  proportions  in  water,  forming 
hydrofluoric  acid.  Its  fumes  are  exceedingly  irritating  to 
the  respiratory  organs,  and  several  chemists  have  lost 
their  lives  by  accidentally  breathing  them. 

Hydrofluoric  acid.  Hydrofluoric  acid,  like  other  acids, 
readily  acts  on  bases  and  metallic  oxides  and  forms  the  cor- 
responding fluorides. .  It  acts  very  vigorously  upon  organic 
matter,  a  single  drop  of  the  concentrated  acid  making  a 
sore  on  the  skin  which  is  slow  in  healing  and  very  painful. 
Its  most  characteristic  property  is  its  action  upon  silicon 
dioxide  (SiO2),  with  which  it  forms  water  and  the  gas 
silicon  tetrafluoride  (SiF4),  as  shown  in  the  equation 

Si02  +  2  HaF2  — +  SiF4  +  2  H2O 


FIG.  104.  A  metal 
U-tube  for  the  prep- 
aration of  fluorine 


268    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Glass  consists  of  certain  compounds  of  silicon  which  are 
likewise  acted  on  by  the  acid,  so  that  it  cannot  be  kept 
in  glass  bottles.  It  is  preserved  in  flasks  made  of  a  wax 
derived  from  petroleum  and  known  as  ceresin  (Fig.  105). 
Ordinary  commercial  hydrofluoric  acid  contains  about 
50  per  cent  of  hydrogen  fluoride. 

Etching.  The  action  of  hydro- 
fluoric acid  on  silicon  compounds  is 
taken  advantage  of  to  etch  designs 
upon  glass.  The  glass  is  painted  over 
with  a  protective  paint  upon  which 
the  acid  will  not  act,  the  parts  which 
it  is  desired  to  make  opaque  being 
left  unprotected.  A  mixture  of  fluor- 
ite  and  sulfuric  acid  is  then  painted 
over  the  vessel,  and  after  a  few  min- 
utes is  washed  off.  Wherever  the 
hydrofluoric  acid  comes  in  contact 
with  the  glass  it  acts  upon  it,  destroy- 
ing its  luster  and  making  it  opaque,  so 
that  the  exposed  design  will  be  etched 
upon  the  clear  glass.  Frosted  glass 
globes  are  often  made  in  this  way, 
but  more  frequently  by  a  sand  blast. 
The  etching  may  also  be  accomplished  by  covering  the  glass 
with  a  thin  layer  of  paraffin,  cutting  the  design  through  the 
wax,  and  then  exposing  the  glass  to  the  fumes  of  the  gas. 


FIG.  105.    A  bottle  made 

out  of  ceresin,  for  holding 

hydrofluoric  acid 


BllOMLNE 

History  and  occurrence.  Bromine  was  discovered  in 
1826  by  the  French  chemist  Ballard,  who  isolated  it 
from  sea  salt.  It  occurs  almost  entirely  in  the  form  of 
sodium  bromide  (NaBr)  and  magnesium  bromide  (MgBr2), 
which  are  found  in  many  springs  and  salt  deposits. 


THE  CHLORINE  FAMILY 


269 


The  Stassfurt  deposits  in  Germany  and  the  salt  waters 
of  Michigan  and  Ohio  are  especially  rich  in  bromides. 

Preparation  of  bromine.  The  laboratory  and  the  com- 
mercial method  for  preparing  bromine  are  as  follows: 

1.  Laboratory  method.  Just  as  chlorine  is  liberated  by 
the  action  of  hydrochloric  acid  on  manganese  dioxide 
(p.  159),  so  bromine  may  be  liberated  by  a  similar  reaction, 
by  using  hydrobromic  acid  in  place  of  hydrochloric : 

4  HBr  +  MnO2 >-  MnBr2  +  Br2  +  2  H2O 

Since  hydrobromic  acid  is  unstable,  it  is  more  con- 
venient to  generate  it  in  the  course  of  the  reaction  by 
using  a  mixture  of  sodium  bromide  and  sulfuric  acid. 
The  equation  for  the  complete  reaction  is  as  follows: 

2NaBr+2H2SO4+MnO2 

The  materials  are  placed 
in  a  retort  A,  arranged  as 
shown  in  Fig.  106.  The 
end  of  the  retort  just 
touches  the  water  in  the 
flask  B,  which  is  partially 
immersed  in  ice  water.  As 
the  contents  of  the  retort 
are  heated  the  bromine 
distills  over,  and  is  col- 
lected in  the  cold  receiver. 


FIG.  106.    An  apparatus  for  the  prepa- 
ration of  bromine  in  the  laboratory 


2.  Commercial  method. 
In  the  United  States 
bromine  is  obtained  commercially  from  salt  water;  the 
bromine  is  liberated  by  electrolysis.  Some  chlorine  is  also 
set  free  in  the  process,  but  this  reacts  with  the  bromides 
present  in  the  water ;  thus, 


2NaBr  +  Cl 


2NaCl 


270    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Properties.  Bromine  is  a  dark-red  liquid  whose  density 
is  3.102.  Its  vapor  has  an  offensive  odor  and  is  very 
irritating  to  the  eyes*  and  throat.  The  liquid  boils  at  63° 
and  solidifies  at  —  7.3°,  but  even  at  ordinary  temperatures 
it  has  a  high  vapor  pressure,  forming  a  reddish-brown  gas 
very  similar  to  nitrogen  dioxide  in  appearance.  100  vol- 
umes of  water  dissolves  about  1  volume  of  bromine  at  20°, 
forming  a  reddish  solution  called  bromine  water.  Bromine 
is  much  more  readily  soluble  in  carbon  disulfide. 

Chemical  conduct  and  uses.  In  chemical  conduct  bromine 
is  very  similar  to  chlorine,  except  that  it  is  less  active. 
It  combines  directly  with  many  of  the  same  elements  with 
which  chlorine  unites,  but  with  less  energy.  It  combines 
with  hydrogen,  and  even  abstracts  it  from  some  of  its 
compounds.  As  would  be  expected,  its  bleaching  action 
is  much  less  marked  than  that  of  chlorine.  Its  solution 
in  water  is  often  used  as  an  oxidizing  agent. 

Bromine  is  used  chiefly  in  the  preparation  of  bromides, 
which  are  employed  to  a  considerable  extent  in  photogra- 
phy and  as  medicinal  agents.  It  is  likewise  used  in  the 
preparation  of  a  number  of  organic  drugs  and  dyestuffs. 

Hydrogen  bromide  (HBr).  When  sulfuric  acid  acts  upon 
a  bromide,  hydrogen  bromide  is  set  free : 

2  NaBr  +  H2SO4 >•  Na2SO4  +  2  HBr 

At  the  same  time  some  bromine  is  liberated,  as  may  be 
seen  from  the  red  fumes  which  appear  and  from  the  odor. 
The  explanation  of  this  is  found  in  the  fact  that  hydrogen 
bromide  is  much  less  stable  than  hydrogen  chloride  and  is 
therefore  more  easily  oxidized.  Concentrated  sulfuric  acid 
is  a  good  oxidizing  agent  (p.  247)  and  oxidizes  a  part  of 
the  hydrogen  bromide,  liberating  bromine: 

H3SO4  +  2  HBr  — >-  H2O  +  H2SO3  +  Br3 


THE  CHLOKINE  FAMILY 


271 


The  pure  compound  is  best  prepared  by  the  action  of  water 
upon  phosphorus  tribroraide,  in  which  reaction  hydrogen 
bromide  and  phosphorous  acid,  P(OH)3  (or  H8POg),  are 
formed.  The  reaction  is  made  clearer  by  the  use  of 
structural  formulas : 


OH 
OH 
OH 


The  preparation  of  hydrogen  bromide  is  conducted  as  fol- 
lows :  Red  phosphorus,  together  with  enough  water  to  cover 
it,  is  placed  in  the  flask  A 
(Fig.  107),  and  bromine  is 
put  into  the  dropping  fun- 
nel B.  By  means  of  the 
stopcock  the  bromine  is  al- 
lowed to  flow  drop  by  drop 
into  the  flask,  the  reaction 
taking  place  without  the 
application  of  heat.  The 
phosphorus  and  bromine 
unite  to  form  PBrg,  which 
then  reacts  with  the  water 
as  expressed  in  the  above 
equation.  The  U-tube  C 
contains  glass  beads  which 

have  been  moistened  with  water  and  rubbed  in  red  phos- 
phorus. Any  bromine  escaping  action  in  the  flask  acts  upon 
the  phosphorus  in  the  U-tube.  The  hydrogen  bromide  is  col- 
lected in  D  by  displacement  of  air,  or  an  aqueous  solution  of 
it  may  be  prepared  as  was  done  in  the  case  of  hydrogen 
chloride  (Fig.  78). 

Properties  and  conduct.  Hydrogen  bromide  is  very  similar 
to  hydrogen  chloride  in  its  properties.  It  is  a  colorless  gas 
and  is  very  soluble  in  water,  1  volume  of  water  dissolving 


FIG.  107.   The  preparation  of  hydro- 
gen bromide 


272    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

612  volumes  of  the  gas  under  standard  conditions.  It 
differs  from  hydrogen  chloride  mainly  in  the  fact  that  it 
is  more  readily  oxidized.  Even  the  oxygen  of  the  air 
gradually  acts  upon  it  to  form  water  and  free  bromine. 

The  aqueous  solution  of  hydrogen  bromide  is  known  as 
hydrobromic  acid.  This  acid  has  properties  similar  to  those 
of  hydrochloric  acid,  but  it  is  less  stable.  It  reacts  with 
metals,  their  oxides,  and  their  hydroxides  to  form  bromides. 

IODINE 

History  and  occurrence.  Iodine  is  present  in  sea  water, 
but  in  relatively  small  quantities.  Certain  seaweeds  absorb 
the  iodine  from  the  water,  in  this  way  concentrating  it 
within  their  tissues.  It  was  from  the  ashes  obtained  by 
the  burning  of  seaweed  that  the  French  chemist  Courtois, 
in  1812,  first  isolated  the  element.  Iodine  is  found  in  cer- 
tain sponges,  oysters,  and  fishes,  but  its  chief  source  is  from 
the  deposits  of  Chile  saltpeter  (sodium  nitrate).  It  is 
interesting  to  note  that  sinall  amounts  of  iodine  exist  in 
the  human  body  in  the  thyroid  gland. 

Preparation.  Iodine  may  be  prepared  in  a  number  of 
ways,  the  principal  methods  being  the  following: 

1.  Laboratory  method.    Iodine  can  readily  be  prepared  in 
the  laboratory  from  an  iodide  by  the  method  used  in  pre- 
paring bromine  (Fig.  106),  except  that  sodium  iodide  is 
substituted  for  sodium  bromide.    It  can  also  be  prepared 
by  passing  chlorine  into  a  solution  of  an  iodide : 

Cl2+2NaI >-2NaCl  +  I2 

2.  Commercial  method.     Formerly  iodine   was   obtained 
entirely  from  the  ashes  of  seaweeds.   While  a  small  amount 
of  the  element  is  still  obtained  from  this  source,  by  far  the 


THE  CHLORINE  FAMILY  273 

largest  supply  comes  from  crude  Chile  saltpeter.  It  is  pres- 
ent in  these  deposits  in  the  form  of  sodium  iodate  (NaIO3) 
and  is  liberated  by  the  action  of  the  sulfites  of  sodium : 

2  NaIO8  +  3  Na2SO3  -f  2  NaHSO3  — *  5  Na2SO4  +  H2O  +  I2 

Properties.  Iodine  is  a  purplish-black  shining  solid 
which,  when  sublimed,  crystallizes  in  brilliant  plates.  It 
has  a.  density  of  4.95,  melts  at  113.5°,  and  boils  at  184.4°. 
The  element  has  a  strong,  unpleasant  odor,  although  the 
odor  is  not  so  disagreeable  as  that  of  chlorine  or  bromine. 
Even  at  ordinary  temperatures  it  gives  off  a  beautiful 
violet  vapor,  which  increases  in  amount  as  heat  is  applied. 
It  is  only  slightly  soluble  in  water,  1  part  being  soluble 
in  3750  parts  of  water  at  15°.  It  is  very  soluble  in  a  solu- 
tion of  potassium  iodide  or  of  hydrogen  iodide,  forming  a 
dark-brown  solution.  It  also  dissolves  in  carbon  disulfide, 
forming  a  violet-colored  solution. 

Chemical  conduct  and  uses.    Chemically,  iodine  is  quite 

similar  to  chlorine   and  bromine,  but  is  still  less  active 

than  bromine.    Both  chlorine  and  bromine  displace  it  from 

its  salts:  2KI+Br2 

2KI  + C12 

When  even  minute  traces  of  iodine  are  added  to  thin  starch 
paste  a  very  intense  blue  color  develops,  and  this  reaction 
forms  a  delicate  test  for  iodine.  A  solution  of  iodine  in 
alcohol  is  called  tincture  of  iodine  and  is  extensively  used 
in  medicine.  Iodine  is -also  largely  used  in  the  preparation 
of  iodides  and  of  certain  dyes  and  organic  drugs.  lodoform, 
a  common  antiseptic,  has  the  formula  CHI8. 

Hydrogen  iodide.  This  compound  is  still  less  stable  than 
hydrogen  bromide ;  it  follows,  therefore,  that  it  cannot  be 
prepared  by  the  action  of  sulfuric  acid  upon  an  iodide 


274    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

(p.  270).  It  is  prepared  by  a  method  similar  to  that  used 
in  the  preparation  of  hydrogen  bromide  ;  namely,  by  the 
action  of  phosphorus  triiodide  on  water.  An  aqueous  solu- 
tion of  the  gas  is  easily  prepared  by  passing  hydrogen  sulfide 
into  water  in  which  finely  divided  iodine  is  suspended  : 

HS+I 


The  hydrogen  iodide  dissolves  in  the  water  as  fast  as 
formed,  while  the  sulfur  is  precipitated  and  is  removed  by 
filtration. 

Properties  of  hydrogen  iodide.  Hydrogen  iodide  resembles 
hydrogen  chloride  and  hydrogen  bromide  in  its  physical 
properties,  being  a  strongly  fuming  colorless  gas.  It  is 
4.37  times  as  heavy  as  air.  At  10°  about  450  volumes 
of  the  gas  dissolves  in  1  volume  of  water.  Owing  to  the 
ease  with  which  the  gas  is  decomposed  into  its  elements, 
it  acts  in  many  respects  like  nascent  hydrogen,  being  a 
strong  reducing  agent.  This  might  be  expected  from  the 
fact  that  it  is  an  endothermic  compound,  as  shown  in 
the  equation 

H2  +  1-2  —  >•  2  HI  -  12,072  cal. 

The  solution  of  hydrogen  iodide  in  water  has  strong  acid 
properties  and  is  known  as  hydriodic  acid. 

Chemical  conduct  of  hydriodic  acid.  Hydriodic  acid  differs 
from  hydrochloric  acid  and  hydrobromic  acid  mainly  in 
the  ease  with  which  it  is  oxidized.  The  freshly  prepared 
solution  is  colorless,  but  soon  turns  brown,  owing  to  the 
liberation  of  iodine  by  the  oxygen  of  the  air: 

4HI  +  O2  —  >-2H2O  +  2I2 

As  the  action  continues,  the  iodine  separates  in  crystalline 
form.  The  acid,  as  well  as  hydrogen  iodide,  is  therefore 
a  strong  reducing  agent. 


THE  CHLORINE  FAMILY  275 

Hydriodic  acid  reacts  with  many  of  the  metals,  as 
well  as  with  their  oxides  and  hydroxides,  to  form  the 
corresponding  salts. 

Salts  of  hydrofluoric  acid,  hydrochloric  acid,  hydrobromic 
acid,  and  hydriodic  acid:  fluorides,  chlorides,  bromides,  iodides. 
Many  of  these  salts  are  well-known  compounds  and  have 
important  uses.  They  can  be  prepared  by  the  usual  methods 
for  preparing  salts ;  namely,  by  the  action  of  the  acid  upon 
the  metals  directly  or  upon  their  oxides  or  hydroxides. 
The  most  important  fluoride  is  the  well-known  calcium 
fluoride  (CaF2),  or  fluorite.  The  chlorides,  bromides,  and 
iodides  resemble  each  other  closely  in  their  properties. 
They  are  all  soluble  in  water  except  the  silver,  lead,  and 
mercurous  salts.  Sodium  chloride  is  the  most  important 
of  the  chlorides.  Potassium  bromide  and  potassium  iodide 
are  used  in  medicine,  while  silver  bromide  and  iodide  are 
largely  used  in  photography. 

THE  OXYGEN  COMPOUNDS  OF  THE  HALOGENS 

The  halogens  have  but  little  affinity  for  oxygen,  only 
chlorine  and  iodine  forming  oxides.  While  several  oxygen 
acids  are  known,  with  few  exceptions  these  are  unstable 
and  exist  only  in  dilute  solution.  Both  the  acids  and  their 
salts  readily  give  up  oxygen.  They  are  therefore  good 
oxidizing  agents  and  are  used  chiefly  for  this  purpose. 

The  oxides.  Chlorine  forms  three  oxides;  namely,  C12O, 
CLjO^  and  C1O2.  They  are  difficult  to  prepare  and  very 
unstable.  Iodine  forms  two  oxides,  of  the  formulas  I2O4 
and  I2O6.  The  latter,  known  as  iodine  pentoxide,  is  the 
best-known  oxide  of  the  group.  It  is  a  white  solid  and 
is  fairly  stable,  although  it  is  decomposed  into  iodine  and 
oxygen  when  heated. 


276    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  oxygen  acids.  The  most  important  of  the  oxygen 
acids  of  the  halogens  are  the  following: 

1.  Hypochlorous  acid  (HCIO).    This  acid  is  unstable  and 
has    been    obtained    only   in   dilute    solutions.    Its   salts, 
known  as  the  hypochlorites,  may  be  prepared  by  passing 
chlorine  into  the  cold  solutions  of  the  hydroxides  of  the 
metals.    Thus,  potassium  hypochlorite  (KC1O)  is  formed 
as  indicated  in  the  following  equation : 

2  KOH  +  C12  — >•  KC10  +  KC1  +  H2O 

Both  the  acid  and  its  salts  easily  give  up  their  oxygen 
and  are  therefore  good  oxidizing  agents. 

2.  Chloric  acid  (HCIOJ.    This  acid  is  more  stable  than 
hypochlorous  acid,  but  has  not  been  obtained  in  a  pure 
state.    Its  salts,  the  chlorates,  may  be  prepared  by  passing 
chlorine  into  a  hot  solution  of  the  hydroxides  of  the  metals. 
Potassium  chlorate,  for  example,  is  easily  obtained  by  pass- 
ing chlorine  into  a  hot  concentrated  solution  of  potassium 
hydroxide : 

3  C12  +  6  KOH  — >-  KC103  +  5  KC1  +  3  H2O 

The  chlorates  are  excellent  oxidizing  agents.  Sodium 
chlorate  and  potassium  chlorate  are  chiefly  used  in  the 
preparation  of  explosives,  fireworks,  and  oxygen. 

Preparation  of  hypochlorites  and  chlorates  by  electrolytic 
methods.  It  will  be  recalled  that  the  electrolysis  of  solutions 
of  potassium  chloride  or  of  sodium  chloride  results  in  the  for- 
mation of  chlorine,  together  with  the  corresponding  hydroxides 
of  metals.  It  is  possible  so  to  regulate  this  process  that  the 
chlorine,  instead  of  being  evolved,  is  retained  in  the  solution, 
together  with  the  hydroxides,  with  which  it  acts  to  form 
hypochlorites  or  chlorates,  according  to  the  equations  given 
above.  This  method  is  now  coming  into  general  use  for  the 


THE  CHLOKIKE  FAMILY  277 

preparation  of  these  salts.  It  is  possible  to  obtain  either 
the  hypochlorites  or  the  chlorates  by  properly  choosing  the 
conditions  of  the  electrolysis. 

3.  Perchloric  acid  (HCIO^.  This  acid  is  likewise  unstable, 
although  it  is  possible  to  obtain  it  in  a  pure  state.  It  is 
a  colorless  liquid,  and  sometimes  decomposes  with  great 
violence.  Potassium  perchlorate,  the  best  known  of  its 
salts,  is  a  white  solid  and,  like  the  salts  of  other  oxygen 
acids  of  chlorine,  is  an  excellent  oxidizing  agent. 

Hypobromous  acid  (HBrO)  and  bromic  acid  (HBr08),  the 
analogues  of  hypochlorous  acid  and  chloric  acid  have  also  been 
prepared  in  dilute  solution.  The  acids,  as  well  as  their  salts,  are 
similar  to  the  corresponding  chlorine  compounds.  lodic  acid 
(HI03)  and  periodic  acid  (HgI06)  are  also  known.  They  are 
both  white  solids.  The  acids,  as  well  as  their  salts,  are  strong 
oxidizing  agents. 

EXERCISES 

1.  How  do  we  account  for  the  fact  that  liquid  hydrogen  fluoride 
does  not  conduct  the  electric  current  ? 

2.  Why  is  sulfuric  acid  used  for  liberating  hydrogen  fluoride 
from  fluorite? 

3.  Why  is   the    formula    for   hydrogen   fluoride   written    H2F2, 
while  that  of  hydrogen  chloride  is  written  HC1? 

4.  What  discoveries    made  by   Moissan  have  been  noted  other 
than  the  preparation  of  fluorine  ? 

5.  C12O  is  the  anhydride  of  what  acid  ? 

6.  A  solution  of  hydrogen  iodide  on  standing  turns  brown.    How 
is  this  accounted  for? 

7.  How  can    bromine    vapor    and    nitrogen    dioxide    be    distin- 
guished from  each  other? 

8.  Write  the  equations  for  the  reaction  which  takes  place  when 
hydrogen  iodide  is  prepared  from  iodine,  phosphorus,  and  water. 

9.  From  their  behavior  toward  sulfuric  acid,  to  what  class  of 
agents  do  hydrobromic  acid  and  hydriodic  acid  belong? 


278    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

10.  Give  the  derivation  of  the  names  of  the  elements  of  the 
chlorine  family. 

11.  Write  the  names  and  formulas  for  the  binary  acids  of  the 
chlorine  group  in  the  order  of  the  stability  of  the  acids. 

12.  What  is  formed  when  a  metal  dissolves  in  each  of  the  follow- 
ing :   nitric  acid ;   dilute  sulfuric  acid ;  concentrated  sulf uric  acid ; 
hydrochloric  acid ;  aqua  regia  ? 

13.  How  'could  you  distinguish  between  a  chloride,  a  bromide, 
and  an  iodide  ? 

14.  In  what  respects  are  the  elements  included  in  the  chlorine 
family  similar? 

15.  What  weight  of  sodium  chloride  is  necessary  to  prepare  suffi- 
cient hydrogen  chloride  to  saturate  1 1.  of  water  under  standard 
conditions  ? 

16.  What  weight  of  fluorite  is  necessary  for  the  preparation  of 
1  kg.  of  the  commercial  hydrofluoric  acid  ? 

17.  The  concentrated  hydrochloric  acid  of  commerce  has  a  den- 
sity of  1.20  and  contains  40  per  cent  of  hydrogen  chloride.    What 
weights  of  salt  and  sulfuric  acid  are  necessary  to  prepare  100  kg.  of 
this  acid? 

18.  What  weights  of  chlorine  and  potassium  hydroxide  are  nec- 
essary for  the  preparation  of  1  kg.  of  potassium  chlorate  ?    Does  this 
process  appeal  to  you  as  an  economical  method  of  preparation  ? 

19.  What  substances  so  far  studied  are  used  as  bleaching  agents  ? 
To  what  is  the  bleaching  action  of  each  due  ? 

20.  On   decomposition  100  1.  of  hydrogen  chloride  would  yield 
how  many  liters  of  hydrogen  and  of  chlorine  respectively? 


CHAPTER  XXIII 
MOLECULAR  WEIGHTS;   ATOMIC  WEIGHTS 

Introduction.  In  early  chapters  two  problems  were  left 
unsolved  for  the  reason  that  their  solution  requires  a  wider 
acquaintance  with  chemical  facts  and  laws  than  had  been 
gained  when  the  problems  were  suggested. 

1.  Molecular  weights.    In  Chapter  VIII  it  was  shown 
that  from  the  results  of  the  careful  analysis  of  a  com- 
pound it  is  easy  to  calculate  a  formula,  provided  we  have 
a  system  of  atomic   weights   and  provided  we   adopt  the 
simplest  formula  possible.   The  method  described  would  lead 
to  the  formula  HO  for  hydrogen  peroxide,  whereas  the 
true  formula  might  be  H2O2  or  H3O3.    To  decide  between 
these  it  is  necessary  to  know  the  molecular  weight  of  the 
compound.   If  it  is  approximately  1 7,  the  simple  formula  is 
correct.    If  it  is  approximately  34,  the  true  formula  is  H2O2. 

2.  The  selection  of  the  atomic  weight  from  the  combining 
weights.    In  Chapter  VII  we  found  that  it  is  easy  to  deter- 
mine a  combining  weight  for  each  element;  but  that  many 
elements  have  more  than  one  combining  weight,  the  one 
a  multiple  of  the  other.    Thus,  the  analysis  of  water  shows 
that  if  we  take  hydrogen  as  unity,  the  combining  weight 
of  oxygen  is  7.94,  while  the  analysis  of  hydrogen  peroxide 
gives  the  combining  weight  of  oxygen  as  15.88.    It  is 
therefore  a  question  of  which  combining  weight  is  to  be 
selected  as  the  true  atomic  weight.    The  clue  to  the  solu- 
tion of  these  two  problems  is  found  in  a  great  generalization 
known  as  Avogadro's  hypothesis. 

279 


280    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Gay-Lussac's  law  of  volumes.  Before  considering  Avo- 
gadro's  hypothesis,  however,  it  is  essential  for  us  to  have 
in  mind  a  law  first  pointed  out  by  Gay-Lussac  (Fig.  28), 
and  known  as  the  law  of  volumes. 

In  the  discussion  of  the  composition  of  hydrogen  chloride 
it  was  stated  that  1  volume  of  hydrogen  combines  with  1  vol- 
ume of  chlorine  to  form  2  volumes  of  hydrogen  chloride.  With 
bromine  and  iodine  similar  combining  ratios  hold  good. 
These  facts  recall  the  simple  volume  relations  already 
noted  in  the  study  of  the  composition  of  steam  (p.  78) 
and  ammonia  (p.  206).  These  relations  may  be  repre- 
sented graphically  in  the  following  way,  the  rectangles 
representing  equal  volumes: 


In  the  early  part  of  the  past  century  Gay-Lussac  studied 
the  volume  relations  of  many  combining  gases,  and  con- 
cluded that  similar  relations  always  hold.  His  observations 
are  summed  up  in  the  following  generalization,  known  as 
the  law  of  volumes  :  When  two  gases  combine  chemically 
there  is  always  a  simple  ratio  between  their  volumes,  and  also 
between  the  volume  of  either  one  of  them  and  that  of  the 
product,  provided  it  is  a  gas.  By  a  simple  ratio  is  meant, 
of  course,  the  ratio  of  integer  numbers,  as  1:2,  or  2  :  3. 

The  relations  expressed  in  the  law  of  volumes  are  so 
simple  and  so  unexpected  that  we  at  once  feel  that  they 
indicate  a  very  simple  ratio  between  the  number  of  mole- 
cules present  in  equal  volumes  of  gases.  As  early  as  1811 
the  Italian  physicist  Avogadro  (Fig.  108)  suggested  that 


MOLECULAK  WEIGHTS;  ATOMIC  WEIGHTS    281 


the  ratios  become  perfectly  intelligible  if  we  assume  that 
equal  volumes  of  any  two  gases  contain  the  same  number  of 
molecules.  This  generalization  is  known  as  the  hypothesis 
of  Avogadro,  and  it  is  in  com- 
plete accord  with  all  that  we 
have  learned  about  gases  since 
Avogadro's  time. 

Avogadro's  hypothesis,  and 
molecular  weights.  Assuming 
the  truth  of  Avogadro's  hypoth- 
esis, we  have  a  simple  means 
of  deciding  upon  the  relative 
weights  of  the  various  kinds  of 
molecules ;  for  if  equal  volumes 
of  two  gases  contain  the  same 
number  of  molecules,  the  weights 
of  the  two  kinds  of  molecules 
must  be  in  the  same  ratio  as  the 
weights  of  the  two  volumes 
made  up  of  these  molecules. 

For  example,  the  weight  of  a 
liter  of  ammonia  is  0.7708  g.  and 
that  of  a  liter  of  hydrogen  chlo- 
ride is  1.6398  g.  These  values 
will  therefore  indicate  the  rela- 
tive weights  of  the  two  kinds  of 
molecules  if  there  is  the  same 
number  of  each  in  a  liter.  If  we 
adopt  some  one  gas  as  a  standard,  we  can  readily  deter- 
mine the  weights  of  all  gaseous  molecules  relatively  to 
those  of  the  standard  gas.  Thus,  if  we  adopt  ammonia 
as  standard  (unity),  the  molecule  of  hydrogen  chloride  is 
2.14  times  as  heavy  as  the  standard. 


FIG.  108.    Statue  erected   at 

Turin,  Italy,  to  the  memory 

of  Avogadro 


282    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Oxygen  as  standard.  It  will  be  seen  that  the  gas  selected 
as  standard  and  the  volume  chosen  for  comparison  will 
make  no  difference,  since  the  weights  are  all  relative  in  any 
case.  But  since  the  molecules  are  all  made  up  of  atoms,  it 
is  important  that  the  standard  chosen  for  atomic  weights 
should  be  in  accord  with  that  chosen  for  molecular  weights. 
For  many  reasons  oxygen  serves  best  for  atomic  weights,  and 
it  is  also  chosen  for  molecular  weights. 

Relative  weights  of  molecules  of  oxygen  and  of  hydrogen. 
In  Chapter  VII  we  saw  that  8  is  the  smallest  integer  that 
can  be  assigned  as  the  combining  weight  of  oxygen  without 
making  the  combining  weight  of  hydrogen  less  than  1. 

Now  the  weight  of  1  1.  of  oxygen  is  1.429  g.  and  that 
of  1  1.  of  hydrogen  is  0.08987  g.  Making  oxygen  8,  the 
ratio  of  these  two  weights  is  8 :  0.504  or  16 : 1.008,  and 
according  to  Avogadro's  hypothesis  these  numbers  must 
represent  the  iv  eights  of  the  two  kinds  of  molecules.  If  we 
can  now  decide  upon  how  many  atoms  are  in  the  mole- 
cule of  oxygen  and  in  that  of  hydrogen,  we  can  at  once 
decide  whether  to  make  the  standard  of  atomic  weights 
the  atom  of  oxygen  at  16  or  at  8. 

Two  atoms  in  the  molecule  of  oxygen  and  of  hydrogen. 
We  have  seen  that  when  hydrogen  and  chlorine  combine, 
the  ratio  by  volume  is  expressed  in  the  equation 

1  volume  hydrogen  + 1  volume  chlorine 

>-  2  volumes  hydrogen  chloride 

Therefore,  according  to  Avogadro's  hypothesis, 

1  molecule  hydrogen  + 1  molecule  chlorine 

>•  2  molecules  hydrogen  chloride 

But  every  molecule  of  hydrogen  chloride  must  contain  at 
least  1  atom ;   therefore,  since  2    molecules  of   hydrogen 


MOLECULAE  WEIGHTS;  ATOMIC  WEIGHTS    283 

chloride  are  formed  from  1  molecule  of  hydrogen,  each 
molecule  of  hydrogen  must  contain  at  least  2  atoms. 

When  hydrogen  and  oxygen  combine  to  form  steam, 
the  ratio  by  volume  is  expressed  in  the  equation 

2  volumes  hydrogen  + 1  volume  oxygen 

>-  2  volumes  steam 

Therefore,  according  to  Avogadro's  hypothesis, 

2  molecules  hydrogen  +  1  molecule  oxygen 

>-  2  molecules  steam 

But  each  molecule  of  steam  has  at  least  1  atom  of  oxygen, 
and  since  2  molecules  of  steam  are  formed  from  1  mole- 
cule of  oxygen,  the  molecule  of  oxygen  must  contain  at 
least  2  atoms. 

There  are  no  facts  known  that  suggest  that  the  molecule 
of  either  hydrogen  or  oxygen  contains  more  than  two  atoms, 
and  so  we  adopt  the  formulas  H2  and  O2  for  these  gases. 

Oxygen  atom,  16;  oxygen  molecule,  32.  Since  the  weights 
of  equal  volumes  of  oxygen  and  hydrogen  (1 1.  each)  are 
in  the  ratio  of  16 : 1.008,  the  weights  of  the  individual 
molecules  are  in  the  same  ratio.  And  since  the  two  kinds 
of  molecules  each  contain  two  atoms,  the  weights  of  the 
two  kinds  of  atoms  are  also  in  the  ratio  16 : 1.008.  Con- 
sequently, if  we  wish  to  have  the  atomic  weight  of  hydro- 
gen greater  than  unity,  we  must  adopt  the  weight  16  and 
not  8  for  the  atomic  weight  of  oxygen.  Since  the  mole- 
cule of  oxygen  consists  of  2  atoms,  we  must  adopt  the 
weight  32  for  the  oxygen  molecule. 

Molecular  weights  from  weight  of  i  liter.  We  have  now 
devised  a  method  of  determining  how  much  heavier  one 
kind  of  molecule  is  than  another,  and  have  fixed  upon 
the  weight  of  one  standard  molecule  (oxygen),  with  which 


284    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

.  all  others  can  be  compared.  The  determination  of  molecu- 
lar weights  now  becomes  easy.  For  example,  1 1.  of  oxy- 
gen weighs  1.429  g.,  while  1 1.  of  hydrogen  chloride  weighs 
1.6398  g.  The  ratio  between  the  weights  of  the  two  kinds 
of  molecules  is  therefore  1.429 : 1.6398.  To  compare  the 
hydrogen  chloride  molecule  with  oxygen  taken  as  32  we 
need  only  solve  the  proportion:  1.429  : 1.6398  =  32:  ^  The 
molecular  weight  of  hydrogen  chloride  (V)  is  therefore  36.7. 

Gram-molecular  volume  equals  22.4  liters.  Having  adopted 
32  as  the  standard  for  oxygen,  it  is  of  interest  to  find  the 
volume  occupied  by  the  gram-molecular  weight  of  this  gas; 
namely,  32  g.  This  volume  will  evidently  be  32  -f- 1.429, 
or  22 A  I.  If  we  construct  a  vessel  of  exactly  this  content 
and  fill  it  with  oxygen  gas,  it  will  contain  just  enough 
molecules  of  oxygen  to  weigh  32  g.,  which  is  our  standard 
weight  for  oxygen. 

If  now  we  replace  the  oxygen  by  another  gas,  say, 
hydrogen  chloride,  we  shall  have  the  same  number  of 
molecules  present.  The  weight  of  hydrogen  chloride  fill- 
ing the  vessel  is  36.45  g.  But  since  there  is  the  same 
number  of  molecules,  the  values  32  and  36.45  must  repre- 
sent the  relative  weights  of  the  two  kinds  of  molecules. 
In  like  manner,  the  weight  of  22.4  1.  of  any  gas  will  give 
a  number  which  expresses  the  weight  of  a  molecule  of 
that  gas  compared  with  the  molecule  of  oxygen  taken  as 
the  standard.  These  relations  are  illustrated  in  Fig.  109. 
We  therefore  reach  the  following  simple  rule :  The  molecu- 
lar weight  of  any  gas  may  be  found  l)y  determining  the  weight 
in  grams  of  22.4  1.  of  the  gas.  The  volume  22.4  1.  is  called 
the  gram-molecular  volume  of  gases.  Owing  to  the  fact 
that  most  gases  do  not  exactly  conform  to  any  of  the 
gas  laws,  the  weight  of  22.4  1.  of  a  gas  is  not  its  precise 
molecular  weight,  but'  is  very  close  to  it. 


MOLECULAR  WEIGHTS;  ATOMIC  WEIGHTS    285 


Other  methods  of  determining  molecular  weights.  It  will 
be  noticed  that  Avogadro's  hypothesis  gives  us  a  method 
by  which  we  can  determine  the  relative  weights  of  the 
molecules  of  two  gases  because  it  enables  us  to  tell  when  we 
are  dealing  with  an  equal  number  of  the  two  kinds  of  mole- 
cules. If  by  any  other  convenient  means  we  can  get  this 
information,  we  can  make  use  of  the  knowledge  so  gained 
to  determine  the  molecular  weights  of  the  two  substances. 


0232g.  HCl  36.45  g.  H2O  18.016  g.  NHS  17.034  g. 

FIG.  109.   The  weight  of  22.4  1.  of  various  gases 

Raoult's  laws.  Two  laws  have  been  formulated  which 
give  us  just  such  information.  They  are  known  as  Raoult's 
laws  and  can  be  stated  as  follows: 

1.  When  weights  of  substances  which  are  proportional  to 
their  molecular  weights  are  dissolved  in  the  same  weight  of  a 
solvent,  the  rise  of  the  boiling  point  is  the  same  in  each  case. 

2.  When  weights  of  substances  which  are  proportional  to 
their  molecular  weights  are  dissolved  in  the  same  weight  of 
a  solvent,  the  lowering  of  the  freezing  point  is  the  same  in 
each  case. 

By  taking  advantage  of  these  laws  it  is  possible  to  deter- 
mine when  two  solutions  contain  the  same  number  of 
molecules  of  two  dissolved  substances,  and  consequently 
to  determine  the  relative  molecular  weights  of  the  two 
substances. 

Molecular  weights  of  the  elements.  When  we  deter- 
mine the  weight  of  22.4  1.  of  the  various  elementary  gases, 
we  reach  some  interesting  conclusions.  Experiment  shows 


286    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


that  the  molecular  weight  of  many  of  them,  such  as  nitro- 
gen, hydrogen,  chlorine,  and  bromine,  give  values  which  are 
twice  the  atomic  weights,  so  that  in  these  cases  the  molecule 
contains  two  atoms  (p.  91).  In  the  case  of  the  metals,  so 
far  as  their  vapors  have  been  studied,  the  molecular  weight 
and  the  atomic  weight  are  the  same,  so  that  the  molecule 
consists  of  a  single  atom. '  The  molecule  of  ozone  contains 
three  atoms  of  oxygen,  so  that  its  formula  is  Og,  while 
the  molecules  of  phosphorus  and  arsenic  contain  four 
atoms,  giving  the  formulas  P4  and  As4. 

Selection  of  atomic  weights  from  combining  weights.  It 
is  now  easy  to  determine  which  multiple  of  the  combining 
weight  of  the  various  elements  shall  be  adopted  as  the  cor- 
rect atomic  weight — the  second  problem  we  set  out  to  solve 
in  this  chapter.  The  mode  of  procedure  will  be  understood 
most  readily  by  an  example  ;  so  let  us  suppose  that  we 
have  found  the  combining  weight  of  nitrogen  to  be  7.005 
and  that  we  wish  to  decide  whether  this  value  or  some 
simple  multiple,  14.01  or  21.015,  is  the  atomic  weight. 

We  first  determine  the  weight  of  22.4  1.  of  a  number  of 
gaseous  compounds  which  we  know  to  contain  nitrogen. 
These  values  are  given  in  the  first  column  of  the  table. 


NAME  OF  GASEOUS 
COMPOUND 

MOLECULAR 

WEIGHT 

(22.4  L.1 

PERCENTAGE  OF 
NITROGEN  BY 
EXPERIMENT 

PART  OF  MOLECU- 
LAR WEIGHT  DUE 
TO  NITROGEN 

Nitrogen  gas     .     .     . 

27.95 

100.00 

27.95 

Nitrous  oxide    .     .     . 

44.13 

63.70 

28.11 

Nitric  oxide  .... 

30.00 

46.74 

14.02 

Ammonia     .... 

17.05 

82.28 

14.03 

Nitric  acid    .... 

63.75 

22.27 

14.30 

We  next  make  a  careful  analysis  of  each  of  these  com- 
pounds to  ascertain  the  percentage  of  nitrogen  present, 


MOLECULAE  WEIGHTS;  ATOMIC  WEIGHTS  287 

placing  the  values  obtained  in  the  second  column.  If  we 
multiply  the  molecular  weight  of  each  compound  by  the 
percentage  of  nitrogen,  the  product  will  be  the  portion  of 
the  molecular  weight  due  to  nitrogen.  But  since  the 
molecules  are  made  up  of  atoms,  the  part  of  a  molecule 
due  to  nitrogen  must  represent  the  sum  of  the  weights  of  the 
nitrogen  atoms  present.  We  notice  that  the  numbers  in  the 
last  column  are  either  very  near  to  14  or  to  twice  14,  and 
that  none  are  near  7.  If  we  examine  a  large  number  of 
nitrogen  compounds,  it  is  reasonable  to  expect  that  we 
shall  find  some  containing  only  one  atom,  and  since  we 
find  none  which  give  a  value  of  less  than  14,  we  assume 
that  this,  and  not  7  or  21  or  28,  represents  the  weight  of 
a  nitrogen  atom. 

Accurate  determination  of  atomic  weights.  The  weight 
of  a  given  volume  of  a  gas  is  difficult  to  determine  with 
great  precision,  and  in  consequence  the  molecular  weights 
of  gases  as  determined  by  experiment  are  usually  subject 
to  a  very  considerable  error.  The  portion  of  nitrogen  in 
22.4  1.  of  the  various  gases  is  therefore  a  little  uncertain, 
as  will  be  seen  from  the  values  in  the  last  column.  All 
these  figures  tell  us  is  that  the  true  value  is  very  near  14. 
The  combining  weight  can  be  very  accurately  determined 
by  the  analysis  of  any  of  these  compounds,  and  is  found 
to  be  7.005.  It  is  therefore  evident  that  the  accurate 
atomic  weight  is  twice  this  value ;  namely,  14.01. 

Summary.  These,  then,  are  the  steps  which  must  be 
taken  to  establish  the  atomic  weight  of  an  element. 

1.  Determine  the  combining  weight  accurately  by  analysis. 

2.  Determine  the  weight  of  22.4  1.  of  a  large  number  of 
gaseous  compounds  of  the  element,  and,  by  analysis,  the  part 
of  the  molecular  weights  due  to  the  element.   The  smallest 
number  so  obtained  will  be  the  approximate  atomic  weight. 


288    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

3.  Multiply  the  combining  weight  by  the  integer  (1,  2, 
or  3)  which  will  give  a  number  close  to  the  approximate 
atomic  weight.  The  number  so  obtained  will  be  the  precise 
atomic  weight. 

Equations  and  volumes  of  gases.  If  we  have  an  equation 
which  expresses  a  reaction  in  which  gaseous  molecules  take 
part,  we  may  make  use  of  Avogadro's  hypothesis  to  predict 
the  volume  changes  which  will  accompany  the  reaction. 
For  example,  the  equation 


states  that  1  gram-molecular  weight  of  hydrogen  combines 
with  1  gram-molecular  weight  of  chlorine  to  give  2  gram- 
molecular  weights  of  hydrogen  chloride.  Now  all  of  these 
substances  are  gases,  and  a  gram-molecular  weight  of  every 
gas  occupies  the  same  volume  ;  namely,  22.4  1.  Conse- 
quently 1  volume  of  hydrogen  will  combine  with  1  volume 
of  chlorine  to  give  2  volumes  of  hydrogen  chloride,  and 
there  will  be  no  change  in  the  volume  due  to  the  reaction 
(save  as  occasioned  by  the  heat  given  off).  The  coefficients 
of  the  molecules  therefore  indicate  the  proportion  by  volume  in 
which  gases  take  part  in  reactions. 

Weight  of  a  liter  of  a  gas.  We  have  found  that  a  gram- 
molecular  weight  of  any  gas  occupies  22.4  1.  If  we  know 
the  molecular  weight  of  a  gas,  we  can  at  once  deduce 
the  weight  of  a  liter  of  the  gas.  For  example,  the  molec- 
ular weight  of  acetylene  (C2H2)  is  26.016.  This  means 
that  26.016  g.  occupies  22.4  1.  Consequently  1  1.  will  weigh 
26.016  -*-  22.4  =  1.1614  g.  In  general,  to  find  the  weight 
of  a  liter  of  any  gas,  divide  its  molecular  weight  by 
22.4.  The  value  so  obtained  will  be  close  enough  to  the 
experimental  value  for  all  practical  purposes. 


NAME                    WEIGHT  OF  1  L. 

COMPOSITION  BY  PEBCI 

Hydrogen  sulfide 
Sulfur  dioxide 

1.5392g. 
2.9266  g. 

S  =  94.11 
S  =  50.05 

H  =  5.89 
O  =  49.95 

Sulfur  trioxide 
Sulfur  chloride 
Sulfuryl  chloride 
Carbon  disulfide 

3.571  g. 
6.027  g. 
6.030  g. 
3.3928  g. 

S  =  40.05 
S  =  47.48 
S  =  23.75 
S  =  84.24 

O  =  59.95 
Cl  =  52.52 
Cl  =  52.53 
C  =  15.76 

MOLECULAR  WEIGHTS;  ATOMIC  WEIGHTS    289 

EXERCISES 

1.  From  the  following  data  calculate  the  atomic  weight  of  sul- 
fur. The  combining  weight  obtained  by  an  analysis  of  sulfur  dioxide 
is  8.015.  The  weight  of  1 1.  of  gas  and  the  compositions  of  a  number 
of  compounds  containing  sulfur  are  as  follows : 


O  =  23.70 


2.  Calculate  the  formulas  for  compounds  of  the  following  per- 
centage compositions: 

MOLECULAR 
WEIGHT 

(1)  8  =  39.07%  O  =  58.49%  H  =  2.44%  81.0 

(2)  Ca  =  29.40%  S  =  23.56%  O  =  47.04%  136.2 

(3)  K  =  38.67%  N  =  13.88%  O  =  47.45%  101.2 

3.  The  molecular  weight  of  ammonia  is  17.06  ;  of  sulfur  dioxide 
is  64.06  ;  of  chlorine  is  70.9.    From  the  molecular  weight  calculate 
the  weight  of  1 1.  of  each  of  these  gases.    Compare  your  results  with 
the  table  in  the  Appendix. 

4.  What  are  the  relative  weights  of  the  molecules  of  hydrogen 
and  hydrogen  chloride,  as  deduced  from  a  weight  of  1 1.  of  each  of 
these  gases? 

5.  Natural  gas  is  largely  composed  of  marsh  gas  (CH4).    When 
this  burns,  the  equation  for  the  reaction  is  as  follows : 

CH4  +  2  O2 >-  CO2  +  2  H2O 

In  burning  100  cu.  ft.  of  this  gas  what  volume  of  oxygen  is  consumed  ? 
What  is  the  volume  of  the  carbon  dioxide  formed  ? 

6.  Why  write  2  O2  rather  than  4  O  in  problem  5  ? 

7.  Determine   the    molecular   weight   of    chloroform    from    the 
following  data :  In  an  experiment  0.2  g.  of  the  liquid  gave  a  volume 
of  42.4  cc.  of  gas  collected  over  water  at  20°  and  740  mm. 

8.  The  percentage  composition  of  chloroform  as  determined  by 
analysis  is  as  follows:  Cl  =  89.11%;  C  =  10.05%;  H  =  0.84%.   From 


290    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

these  figures  and  taking  into  consideration  the  molecular  weight  as 
determined  in  problem  7,  calculate  the  formula  of  chloroform. 

9.  Solve  problem  20,  Chapter  XXII,  without  using  molecular 
weights.   Compare  your  results. 

10.  When  ammonia  is  heated  to  ignition  in  an  atmosphere  of 
oxygen  it  burns,  forming  water  and  liberating  nitrogen.    What  is 
the  relation  between  the  volume  of  the  ammonia  burned  and  that  of 
the  oxygen  required  for  its  combustion?   between  the  volume  of 
the  ammonia  burned  and  the  nitrogen  liberated  in  its  combustion? 

11.  Compare  the  volume  of  hydrogen  sulfide  with  that  of  the 
oxygen  required  for  its  complete  combustion. 


CHAPTER  XXIV 

CARBON  MONOXIDE ;   CARBONIC  ACID ;   HYDROCARBONS 
I 

Introductory.  In  connection  with  the  occurrence  of 
carbon  (p.  116),  attention  was  called  to  the  fact  that 
this  element  is  widely  distributed'  in  nature  and  that  the 
number  of  its  compounds  is  very  large.  Over  200,000  of 
them  have  been  described,  and  many  newly  discovered 
ones  are  constantly  being  added  to  the  list.  Because  the 
number  of  carbon  compounds  is  so  great  and  also  because 
certain  characteristics  distinguish  them  from  the  compounds 
of  other  elements,  it  has  been  found  convenient  to  group 
them  under  the  general  heading  of  organic  chemistry  and 
to  postpone  their  study  until  the  introductory  course  in 
chemistry  has  been  concluded.  Nevertheless,  even  an  in- 
troductory course  must  include  a  limited  number  of  the 
more  common  compounds  of  carbon. 

THE  OXIDES  OF  CARBON  AND  CARBONIC  ACID 

Carbon  forms  three  oxides ;  namely,  carbon  monoxide 
(CO),  carbon  dioxide  (CO2),  carbon  suboxide  (C3O2). 
They  are  all  colorless  gases.  But  little  is  known  of  the 
suboxide,  and  no  further  mention  will  be  made  of  it. 
Carbon  dioxide  has  been  described  in  connection  with 
carbon  (p.  122). 

Carbon  monoxide  (CO).  Carbon  monoxide  occurs  in  the 
gases  issuing  from  volcanoes.  It  can  be  prepared  in  a  num- 
ber of  ways,  the  most  important  of  which  are  the  following : 
291 


292    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

1.  By  the  partial  reduction  of  carbon  dioxide.   When  car- 
bon dioxide  is  conducted  over  highly  heated  carbon  the 
monoxide  results: 

CO2  +  C  —  >-2CO 

When  coal  burns  in  a  stove  carbon  dioxide  is  at  first 
formed  in  the  free  supply  of  air,  but  as  the  hot  gas  rises 
through  the  glowing  coals  it  is  reduced  to  carbon  mon- 
oxide. When  this  gas  comes  in  contact  with  the  air 
above  the  coal,  the  gas  combines  with  oxygen  to  form  car- 
bon dioxide,  burning  with  the  blue  flame  so  often  noticed 
above  a  bed  of  coals,  especially  in  the  case  of  hard  coal. 

2.  By  the  decomposition  of  formic  acid.    In  the  laboratory 
carbon   monoxide  is  usually  prepared  by  heating  formic 
acid  (CH2O2)  or  its  sodium  salt  with  sulfuric  acid  : 

CH202  —  ^H20  +  C0 

The  sulfuric  acid  assists  in  the  reaction  by  combining 
with  the  water  formed.  The  carbon  monoxide  may  be 
collected  over  water,  since  it  is  but  slightly  soluble. 


Oxalic  acid  (G^H.^0^)  may  be  used  in  place  of  formic  acid  : 
C2H204  -  >•  CO  +  C02  +  H20 

In  this  case,  however,  it  is  necessary  to  pass  the  gaseous 
mixture  through  a  solution  of  sodium  hydroxide  to  remove 
the  carbon  dioxide. 

Properties.  Carbon  monoxide  is  a  colorless  and  odorless 
gas.  It  is  0.967  times  as  heavy  as  air  and  is  very  diffi- 
cult to  liquefy.  It  burns  in  air  or  in  oxygen  with  a  blue 
flame,  forming  carbon  dioxide.  It  combines  with  chlorine 
to  form  the  gas  COC12,  known  as  phosgene,  or  carbonyl 
chloride.  It  also  combines  directly  with  some  of  the 
metals,  as  nickel  and  iron.  Because  of  its  affinity  for 


CABBON  MONOXIDE 


293 


oxygen  it  is  a  good  reducing  agent.  For  example,  when 
it  is  passed  over  hot  copper  oxide,  the  oxygen  is  with- 
drawn according  to  the  following  equation : 

CuO  +  CO *•  Cu  +  CO2 

Carbon  monoxide  is  very  poisonous,  and  being  nearly  odor- 
less, it  is  a  very  treacherous  poison.  Deaths  not  infrequently 
result  from  the  stoppage  of  stovepipes  or  chimneys.  The  draft 
of  air  is  diminished  to  such  an  extent  that  carbon  monoxide 


FIG.  110.   The  reducing    power  of  carbon  monoxide,  as  shown  in  the 
reduction  of  copper  oxide  by  the  hot  gas 

rather  than  dioxide  is  the  main  product  of  combustion  and, 
not  having  egress  through  the  chimney,  escapes  into  the  room. 
It  is  an  interesting  fact  that  birds  are  very  sensitive  to  this 
gas.  In  mine  explosions  carbon  monoxide  is  always  formed, 
and  rescuers  often  carry  canaries  with  them,  the  death  of  the 
birds  warning  the  rescuers  of  their  own  peril. 

The  reducing  power  of  carbon  monoxide.  Fig.  110  illustrates  a 
method  of  showing  the  reducing  power  of  carbon  monoxide. 
The  gas  is  generated  by  gently  heating  a  mixture  of  formic 
acid  and  sulfuric  acid  in  the  flask  A.  The  bottle  B  contains  a 
little  water  to  wash  the  gas.  C  is  a  hard-glass  tube  containing 
copper  oxide,  which  is  heated  by  a  burner.  The  black  copper 


294    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

oxide  is  reduced  to  reddish  metallic  copper  by  the  carbon 
monoxide,  which  is  in  turn  changed  to  carbon  dioxide.  The 
formation  of  the  carbon  dioxide  is  shown  by  the  precipitate  in 
the  calcium  hydroxide  solution  in  D.  Any  unchanged  carbon 
monoxide  is  collected  over  water  in  F.  If  oxalic  acid  is  used 
in  preparing  the  oxide,  bottle  E  should  contain  a  solution  of 
sodium  hydroxide  to  remove  the  carbon  dioxide  which  is 
formed  along  with  the  monoxide. 

Carbonic  acid  (H2C03).  Like  most  of  the  oxides  of  the 
nonmetallic  elements,  carbon  dioxide  is  an  acid  anhydride. 
It  combines  with  water  to  form  an  acid  of  the  formula 
H2CO3,  called  carbonic  acid  : 

H20  +  C02  —  )~H2C03 

The  acid  is,  however,  very  unstable  and  cannot  be  iso- 
lated. Only  a  very  small  amount  of  it  is  actually  formed 
when  carbon  dioxide  is  passed  into  water,  as  is  evident 
from  the  small  solubility  of  the  gas.  If,  however,  a  base 
is  present  in  the  water,  salts  of  carbonic  acid  are  formed, 
and  these  are  quite  stable: 

2  NaOH  +  H2CO3  —  >-  Na2CO3  +  2  H2O 

Action  of  carbon  dioxide  on  bases.  This  conduct  is  explained 
by  the  principles  of  reversible  reactions.  The  equation 


is  reversible,  and  the  extent  to  which  the  reaction  progresses 
depends  upon  the  relative  concentrations  of  each  of  the  three 
factors  in  it.  Equilibrium  is  ordinarily  reached  when  very 
little  H2C08  is  formed.  If  a  base  is  present  in  the  water  to 
combine  with  the  H2C08  as  fast  as  it  is  formed,  all  of  the  C02 
is  converted  into  H2C08  and  thence  into  a  carbonate. 

Salts  of  carbonic  acid  ;  carbonates.  The  carbonates  form 
an  important  class  of  salts.  Limestone,  shells,  and  marble 
are  largely  calcium  carbonate  (CaCO8),  common  washing 


CARBONIC  ACID  295 

soda  is  sodium  carbonate  (Na2COg),  and  baking  soda  is 
sodium  acid  carbonate  (NaHCO8).  The  carbonates  of 
sodium,  potassium,  and  ammonium  only  are  soluble,  and 
these  can  be  made  by  the  action  of  carbon  dioxide  on 
solutions  of  the  bases,  as  has  just  been  explained. 

The  insoluble  carbonates  are  formed  as  precipitates 
when  soluble  salts  are  treated  with  a  solution  of  a  soluble 
carbonate.  Thus,  the  insoluble  calcium  carbonate  can  be 
made  by  bringing  together  solutions  of  calcium  chloride 
and  sodium  carbonate,  as  follows: 

CaCl2  +  Na2C03  -  >-  CaCO3  +  2  NaCl 

Most  of  the  carbonates  are  decomposed  by  heat,  yielding 
carbon  dioxide  and  an  oxide  of  the  metal.  Thus,  calcium 
oxide  (lime)  is  made  by  heating  calcium  carbonate  (lime- 
stone), as  follows  : 

CaC'O 


The  carbonates  are  readily  acted  upon  by  acids,  liberating 
carbon  dioxide  (p.  123). 

Action  of  carbon  dioxide  on  calcium  hydroxide.  If  carbon 
dioxide  is  passed  into  a  solution  of  calcium  hydroxide  (lime- 
water),  calcium  carbonate  is  at  first  precipitated  (p.  125): 

H20  +  C02  -  >-H2C08, 
Ca(OH)2  +  H2C08  -  *•  CaC08  +  2  H20 

If  the  current  of  carbon  dioxide  is  continued,  the  precipitate 
soon  dissolves  because  the  excess  of  carbonic  acid  forms  cal- 
cium acid  carbonate,  which  is  soluble  : 

CaC08  +  H3C08  -  »-Ca(HC08)2 

If  now  the  solution  is  heated,  the  acid  carbonate  is  decomposed 
and  calcium  carbonate  once  more  precipitated  : 

Ca(HC08)2  -  **  CaC08  +  HaCO, 


296    AN  ELEMENTARY  STUDY  OF  CHEMISTEY 

Cyanogen  (C2N2)  and  hydrogen  cyanide  (HCN).  At  high 
temperatures  carbon  unites  with  nitrogen  to  form  the 
colorless,  very  poisonous  gas,  cyanogen  (C2N2).  With 
hydrogen  and  nitrogen  it  forms  hydrogen  cyanide  (HCN). 
This  is  a  colorless  liquid  boiling  at  26°.  It  has  a  peculiar 
odor  suggesting  peach  kernels  and  is  extremely  poisonous 
either  when  its  vapor  is  inhaled  or  when  the  liquid  is  taken 
internally.  The  vapor  is  often  used  to  destroy  insects.  It 
is  soluble  in  water  in  all  proportions,  forming  the  solution 
known  as  hydrocyanic  acid  or,  more  frequently,  as  prussic 
acid.  It  is  a  very  weak  acid.  Its  salts  are  called  cyanides 
and,  like  the  acid  itself,  are  very  poisonous.  Sodium  cya- 
nide (NaCN)  and  potassium  cyanide  (KCN)  are  white 
solids.  Their  solutions  readily  dissolve  gold  and  are  often 
used  for  extracting  gold  from  its  ores. 

Since  hydrogen  cyanide  is  very  volatile,  it  is  easily 
liberated  from  the  cyanides  by  the  action  of  sulfuric  acid 
(p.  225),  and  this  is  the  usual  method  for  preparing  it: 

KCN  +  H2SO4  —  >-  HCN  +  KHSO4 

Structural  formula  of  hydrogen  cyanide.  In  some  reactions 
hydrogen  cyanide  acts  as  though  it  had  the  formula  H—  C  =  N, 
while  other  reactions  indicate  the  formula  H—  N—  C.  These 
facts  can  be  explained  by  the  assumption  that  what  we  know 
as  hydrogen  cyanide  is  really  a  mixture  of  two  compounds  in 
equilibrium  with  each  other  as  indicated  in  the  following 
equation  : 


The  hydrocarbons.  Carbon  and  hydrogen  unite  to  form 
a  great  many  compounds.  These  are  known  collectively 
as  the  hydrocarbons.  Their  importance  may  be  inferred 
from  the  fact  that,  mixed  in  varying  proportions,  they  con- 
stitute such  valuable  substances  as  natural  gas,  gasoline, 


HYDROCARBONS  297 

kerosene,  vaseline,  and  paraffin ;  moreover,  from  them  are 
prepared  most  of  our  dyes  and  some  of  our  most  powerful 
explosives. 

Classes  of  hydrocarbons.  In  order  to  simplify  the  study 
of  the  hydrocarbons,  it  is  convenient  to  arrange  them  in 
groups,  or  series.  The  most  important  of  these  are  the 
methane  series  and  the  benzene  series. 

1.  The  methane  series  of  hydrocarbons.    In  the  following 
table  are  given  the  names,  formulas,  and  boiling  points 
of  some  of  the  members  of  the  methane  series: 

NAME  AND  FORMULA  BPOINT°  NAME  AND  FORMULA  B^oiNTG 

Methane  (CH4)      ...  - 160°  Pentane  (C6H12)      .     .     .  +  36° 

Ethane  (C2H6)       .     .     .     -  93°  Hexane  (C6H14)      .     .     .   +  69° 

Propane  (C3H8)     ...     -  45°  Heptane  (C7H16)     .     .     .   +  98° 

Butane  (C4H10)      ...        +1°  General  formula  (CnH2n  +  2) 

Each  member  of  this  series  differs  from  the  one  preceding 
it  by  the  group  of  atoms  (CH2),  and  the  boiling  points 
gradually  increase.  All  the  members  of  this  series  are 
known  up  to  the  one  having  the  formula  C28H5g.  The 
lower  members  are  gases,  the  intermediate  members  are 
liquids,  and  the  higher  members  are  solids.  They  are  all 
combustible. 

The  great  source  of  the  methane  hydrocarbons  is  petro- 
leum, an  oil  obtained  by  boring  deep  wells  into  the  earth 
in  certain  localities  (Fig.  111).  This  oil  is  composed  prin- 
cipally of  liquid  hydrocarbons  in  which  are  dissolved  both 
gaseous  and  solid  hydrocarbons. 

2.  The  benzene  series  of  hydrocarbons.    The  most  impor- 
tant members  of  the  benzene  series  are  benzene  (C6H6) 
and  toluene  (C7Hg).   Both  are  colorless  liquids,  the  former 
boiling  at   80.2°   and  the   latter  at  110°.     The  benzene 
hydrocarbons  are  obtained  from  coal  tar,  a  sticky  black 


298    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

liquid  obtained  in  the  manufacture  of  coke  and  coal  gas. 
They  are  of  great  importance  in  the  manufacture  of  dyes 
and  explosives. 

Production  and  refining  of  petroleum.  The  chief  oil- 
producing  regions  in  the  United  States  are  in  California, 
Oklahoma,  Illinois,  and  Texas.  The  United  States  produces 


FIG.  111.   Typical  scene  in  an  oil  field,  showing  the  wells  and  the 
storage  tanks 

annually  about  300,000,000  barrels,  which  is  about  two 
thirds  of  the  world's  output.  The  oil  is  pumped  to  the 
surface  and  scored  in  large  tanks  (Fig.  Ill)  until  refined. 
Petroleum,  except  when  the  crude  product  is  used  as  a 
fuel,  is  always  subjected  to  a  refining  process  in  which 
the  oil  is  separated  into  different  constituents  which  are 
then  purified.  In  this  process  the  crude  oil  is  run  into 
large  iron  stills  (Fig.  112)  and  subjected  to  distillation. 


HYDROCARBONS 


299 


The  distillates  which  pass  over  between  certain  limits  of 
temperature  are  kept  separate  from  each  other  and  serve 
for  different  purposes.  Thus,  the  liquid  distilling  between 
approximately  70°  and  150°  is  known  as  naphtha,  that  dis- 
tilling between  150°  and  300°  is  ordinary  kerosene  (coal  oil), 
while  the  oils  passing  over  above  300°  are  used  as  lubricat- 
ing oils.  The  oils  that  are  semisolid  at  ordinary  temperature 


FIG.  112.    Stills  for  refining  petroleum 

constitute  vaseline.  The  liquid  remaining  after  the  higher- 
boiling  oils  have  distilled  over  is  chilled,  whereupon  the 
solid  constituents  dissolved  in  the  oil  separate.  These  are 
filtered  off  and  yield  such  products  as  paraffin  and  ceresin. 

The  naphthas.  A  number  of  different  naphthas  are  recog- 
nized commercially,  differing  in  boiling  points  and  densities. 
Those  of  low  boiling  point  are  called  gasoline  and  are  used  in 
gasoline  engines  and  as  a  fuel ;  those  of  a  higher  boiling  point 
are  used  in  making  paints.  Benzine  is  a  high-boiling  naphtha, 
and  being  a  good  solvent  for  such  organic  substances  as  fats 
and  oils,  is  used  in  cleaning  fabrics  (dry-cleaning). 


300   'AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  various  products  obtained  from  petroleum  by  distilla- 
tion are  purified,  usually  by  washing  first  with  sulf uric  acid, 
then  with  sodium  hydroxide,  and  finally  with  water. 

It  is  evident  from  the  method  of  preparation  that  the  prod- 
ucts obtained  from  petroleum  such  as  gasoline  and  kerosene 
are -not  definite  hydrocarbons  but  mixtures  of  hydrocarbons 
that  boil  between  certain  limits  of  temperature.  (One  should 
not  confuse  the  two  products,  benzene  and  benzine :  the  former 
is  a  definite  hydrocarbon,  C6H6,  obtained  chiefly  from  coal  tar, 
while  the  latter  is  a  mixture  of  low-boiling  hydrocarbons 
obtained  from  petroleum.) 

Because  of  the  ease  with  which  benzine  burns,  as  well  as  of 
the  explosive  character  of  a  mixture  of  its  vapor  and  air,  many 
accidents  result  from  its  use,  especially  when  it  is  employed 
in  our  homes  for  cleaning  fabrics.  The  greatest  care  must  be 
taken  in  cleaning  silk,  since  friction  often  causes  a  spark. 

The  cracking  of  oils.  Formerly  kerosene  was  the  most  impor- 
tant of  the  products  obtained  from  petroleum.  At  present, 
however,  gasoline  is  by  far  the  most  valuable,  so  that  every 
effort  is  now  made  to  increase  the  yield  of  gasoline.  To  accom- 
plish this  the  distillation  is  carried  on  under  conditions  that 
tend  to  decompose  the  heavier  molecules  making  up  the  higher- 
boiling  liquids  into  the  simpler  molecules  which  constitute 
liquids  of  lower  boiling  points.  The  process  is  known  as  the 
cracking  of  oils.  It  consists  essentially  in  vaporizing  the  oils 
and  then  heating  the  vapor  under  considerable  pressure.  Good 
results  are  obtained  at  a  temperature  of  from  500°  to  550°  and 
under  a  pressure  of  12  atmospheres.  It  is  of  great  interest  to 
note  that  it  is  possible  in  this  way  not  only  greatly  to  increase 
the  yield  of  gasoline  obtainable  from  a  given  sample  of  petro- 
leum but  also  by  selecting  proper  conditions  of  temperature 
and  pressure,  to  bring  about  reactions  that  result  in  the  forma- 
tion of  certain  hydrocarbons  of  the  benzene  series,  especially 
benzene  and  toluene.  This  method  is  being  used  to  some 
extent  in  the  preparation  of  benzene  and  toluene,  since  both 
of  these  compounds  are  very  valuable  for  the  preparation  of 
dyes  and  explosives. 


HYDROCARBONS 


301 


FIG.  113.  The  miner's 
safety  lamp 


Methane  (marsh  gas)  (CH4).  Methane  is  the  first  member 
of  the  methane  series  of  hydrocarbons  (p.  297).  It  consti- 
tutes from  90  to  95  per  cent  of  natural 
gas.  It  is  formed  in  "marshes  by  the 
decay  of  vegetable  matter  under  water, 
and  bubbles  of  the  gas  are  often  seen 
to  rise  when  the  dead  leaves  on  the 
bottom  of  pools  are  stirred.  It  also  col- 
lects in  mines,  and,  when  mixed  with 
air,  is  called  fire  damp  by  the  miners, 
because  of  its  great  inflammability, 
damp  being  an  old  name  for  gas.  It 
is  formed  when  organic  matter,  such 
as  coal  or  wood,  is  heated  in  closed 
vessels,  and  is  therefore  a  principal 
constituent  of  coal  gas.  Pure  methane 
is  a  colorless,  odorless  gas  about  one 
half  as  heavy  as  air.  It  is  but  slightly  soluble  in  water. 
When  ignited  it  burns  with  a  pale-blue  flame: 

CH4  +  2  O2-^CO2  +  2  H2O  +  213,500  cal. 

Safety  lamp.  Fortunately  the  kin- 
dling temperature  of  fire  damp  is  high, 
and  its  flame  may  be  extinguished  by 
cooling.  In  1815  Sir  Humphry  Davy 
(Fig.  80)  invented  a  miner's  lamp  based 
on  this  principle,  in  which  the  usual 
chimney  of  a  lantern  is  replaced  by  a 
wire  gauze  (Fig.  113).  An  explosion 
flame  starting  at  the  wick  is  so  cooled 
by  the  metal  wire  that  ignition  ceases 
and  the  explosion  is  confined  to  the 
interior  of  the  lamp.  The  principle 
may  be  demonstrated  by  holding  a  wire  gauze  a  few  inches 
above  a  Bunsen  flame  and  parallel  with  the  table  (Fig.  114). 


FIG.  114.  An  experiment 

to  illustrate  the  principle 

of  the  safety  lamp 


302    AN  ELEMENTAEY  STUDY  OF  CHEMISTRY 

When  the  gas  is  turned  on  and  a  light  applied  above  the 
gauze,  the  resulting  flame  rests  upon  the  gauze,  but  does  not 
pass  through  it  to  the  burner. 

Halogen  derivatives  of  methane.  As  a  rule  the  hydrogen  pres- 
ent in  a  hydrocarbon  may  be  displaced  by  a  halogen  element, 
atom  for  atom.  In  this  way  there  are  formed  from  methane 
a  number  of  derivatives,  the  most  important  of  which  are  the 
following : 

Chloroform  (CHClg),  a  heavy  colorless  liquid  boiling  at  61°, 
is  the  well-known  anesthetic  used  in  surgery.  Carbon  tetra- 
chloride  (CC14)  resembles  chloroform  in  appearance.  It  is  a 
good  solvent,  especially  for  fatty  substances.  It  is  often  used 
to  remove  grease  spots  from  fabrics,  and  is  sold  for  this  pur- 
pose under  the  name  of  carbona.  It  possesses  the  advantage 
over  benzine  of  being  noninflammable,  but  is  more  expensive. 
lodoform  (CHI3)  is  a  yellow  crystalline  solid  and  is  largely 
used  as  an  antiseptic  in  the  treatment  of  wounds. 

Acetylene  (C2H2).  This  hydrocarbon  is  a  colorless  gas 
and  is  now  made  in  large  quantities  by  the  action  of 
water  on  calcium  carbide  (CaC0): 

CaC2  +  2  H20 »•  C2H2  +  Ca(OH)2 

The  gas  when  pure  has  a  faint,  pleasant  odor,  the  dis- 
agreeable odor  of  ordinary  acetylene  being  due  to  im- 
purities. It  is  an  endothermic  compound;  that  is,  it 
decomposes  with  evolution  of  heat: 

C2H2  — >•  2  C  +  H2  +  49,300  cal. 

Acetylene,  with  the  proper  admixture  of  air,  burns  with  a 
brilliant  white  light.  The  flame  is  very  hot  because  to 
the  heat  of  combustion  of  the  carbon  and  hydrogen  present 
there  is  added  the  heat  of  decomposition  of  the  acetylene 
undergoing  combustion: 

2  C2H2  +  5  O2  — >•  4  CO2  +  2  H2O  +  2  x  301,630  cal. 


HYDROCARBONS  303 

Acetylene  is  very  explosive  when  subjected  to  pressure. 
It  has  been  found  that  the  gas  can  be  compressed  with 
safety,  however,  by  forcing  it  at  low  temperatures  into 
metal  cylinders  completely  filled  with  some  porous  material 
such  as  a  mixture  of  asbestos  and  cotton,  which  has  been 
partially  saturated  with  certain  liquids  (acetone,  a  liquid 
obtained  by  the  destructive  distillation  of  wood,  is  often 


FIG.  115.   Cutting  an  iron  plate  by  means  of  the  oxyacetyleue  blowpipe 

used).  These  liquids  absorb  large  volumes  of  the  gas, 
and  under  such  conditions  it  is  not  explosive.  Stored  in 
this  way  the  gas  is  now  a  common  article  of  commerce. 

Uses  of  acetylene.  As  an  illuminant,  acetylene  is  often 
used  in  places  where  electric  lights  are  not  available.  The 
chief  use  of  the  gas  at  present  is  in  the  cutting  and  weld- 
ing of  metals  and  in  burning  out  the  carbon  deposited  in 
the  cylinders  of  gasoline  engines.  For  these  purposes 
acetylene  is  burned  in  pure  oxygen  in  a  "form  of  apparatus 


304    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

known  as  the  oxyacetylene  blowpipe,  which  is  almost  exactly 
like  the  oxyhydrogen  blowpipe.  A  temperature  of  about 
2700°  may  be  obtained  in  this  way.  This  blowpipe  has 
been  found  especially  useful  in  dismantling  iron  structures, 
since  the  tip  of  the  flame,  when  drawn  slowly  over  the 
metal,  burns  it  at  the  point  of  contact  (Fig.  115)  and  so 
cuts  the  metal  into  pieces. 


EXERCISES 

1.  Determine  the  percentage  composition  of  carbon  monoxide 
and  carbon  dioxide ;  of  marsh  gas  and  acetylene. 

2.  How  could  you  prove  that  carbonic  acid  is  formed  when 
carbon  dioxide  is  passed  into  water?    What  other  gases  have  we 
studied  that  combine  with  water  ? 

3.  Why  do  most  acids  decompose  carbonates? 

4.  What  compound  would  be  formed  by  passing  carbon  dioxide 
into  a  solution  of  potassium  hydroxide?    Write  the  equation. 

5.  Suggest  a  method  for  the  preparation  of  ammonium  carbonate. 

6.  In   what    respect   are    carbonic    acid    and    sulfurous    acid 
similar  ? 

7.  How    could    you    distinguish    between    sodium    carbonate 
(Na2CO8)  and  sodium  sulfite  (Na2SO8)  ? 

8.  Give  the  reasons  why  the  reaction  which  takes  place  when 
calcium  acid  carbonate  is  heated  completes  itself. 

9.  Could  a   solution  of   sodium   hydroxide  be  substituted  for 
the  solution  of  calcium  hydroxide  in  testing  for  carbon  dioxide  ? 

10.  How  could  you  distinguish  between  gasoline  and  kerosene  ? 

11.  Could  asbestos  fibers  be  used  to  replace  the  wire  in  a  safety 
lamp? 

12.  What  weight  of  formic  acid  is  necessary  for  the  preparation 
of  10  1.  of  carbon  monoxide  ? 

13.  (a)  What  volume  of  oxygen  is  required  for  combustion  of 
10  1.  of  carbon  monoxide  ?    (b)  What  is  the  volume  of  the  carbon 
dioxide  formed  ? 


HYDROCARBONS  305 

14.  What  volume  of  oxygen  would  be  required  to  burn  10  1.  of 
methane  ?  of  acetylene  ?    What  volume  of  carbon  dioxide  would  be 
formed  in  each  case  ? 

15.  What  weight  of  formic  acid  would  be  required  in  the  prepa- 
ration of  sufficient  carbon  monoxide  to  reduce  10  g.  of  copper  oxide 
to  copper? 

16.  What  weight  of  sodium  hydroxide  is  necessary  to  neutralize 
the  carbonic   acid  formed  by  the  action  of  hydrochloric  acid  on 
100  g.  of  calcium  carbonate  ? 

17.  On  the  supposition  that  calcium  carbide  costs  12  cents  a 
kilogram,  what  would  be  the  cost  of  an  amount  sufficient  to  generate 
100  1.  of  acetylene  measured  at  20°  and  740  mm.  ? 

18.  (a)  How  many  calories  of  heat  are  evolved  in  the  combustion 
of  100  1.  of  acetylene  ?    (7>)  What  weight  of  water  at  20°  would  this 
heat  convert  into  steam  at  100°  ? 

19.  Supposing  that  gasoline  is  pure  heptane,  what  volume  of  air 
is  necessary  to  burn  1  kg.  of  the  liquid  ?   What  volume  of  carbon 
dioxide  would  be  formed? 


CHAPTER  XXV 
FUELS;  FLAMES;  ELECTRIC  FURNACES 

Fuels.  Many  substances  are  used  as  sources  of  heat, 
the  most  important  being  the  various  fuel  gases,  together 
with  coal,  wood,  and  petroleum.  The  composition  of  several 
of  these  fuel  gases  is  given  in  the  table  on  page  311. 
Most  of  them  serve  as  illuminants  as  well  as  for  fuels. 

Coal  gas.  It  has  been  known  for  several  centuries  that 
when  soft,  or  bituminous,  coal  is  heated  out  of  contact 
with  air,  combustible  gases  are  formed ;  indeed,  gas  ob- 
tained in  this  way  was  used  for  street  lighting  in  London 
and  Paris  more  than  a  hundred  years  ago. 

The  manufacture  of  coal  gas.  The  manufacture  of  coal  gas  is 
represented  in  a  diagrammatic  way  in  Fig.  116.  The  coal  is 
introduced  into  a  closed  retort  A  and  heated  by  the  fire  below. 
A  number  of  these  retorts  are  placed  in  horizontal  rows,  each 
being  furnished  with  a  delivery  pipe.  This  delivery  pipe  leads 
into  a  large  pipe  B  (known  as  the  hydraulic  main)  which  runs 
at  right  angles  to  the  retort.  The  application  of  heat  causes  the 
coal  to  undergo  complex  changes  which  result  in  the  formation 
of  a  large  number  of  compounds.  These  compounds  escape 
through  the  delivery  pipe  into  the  hydraulic  main.  From  the 
hydraulic  main  the  impure  gas  then  passes  into  a  series  of 
pipes  C,  in  which  it  is  cooled.  Here  is  deposited  a  thick,  tarry 
mass  known  as  coal  tar,  which  is  a  mixture  of  a  large  number 
of  liquid  and  solid  products  formed  in  the  heating  of  the  coal. 
On  the  top  of  the  tar  there  collects  a  liquid,  mostly  water, 
containing  ammonia  and  known  as  the  ammoniacal  liquor.  In 
the  scrubber  D  the  gas  passes  through  a  column  of  loose  coke, 
over  which  water  is  sprayed,  where  it  is  still  further  cooled 
306 


FUELS;  FLAMES;  ELECTRIC  FURNACES     307 

and  to  some  extent  purified  from  soluble  gases,  such  as  hydro- 
gen sulfide  and  ammonia.  In  the  purifier  E  it  passes  over  a 
bed  of  lime  or  of  iron  oxide,  which  removes  the  remainder  of 
the  sulfur  compounds,  and  from  this  it  enters  the  large  gas 
holder  F,  from  which  it  is  distributed  to  consumers. 

The  great  bulk  of  the  carbon  remains  in  the  retort  as  coke 
and  as  retort  carbon.  The  yield  of  gas,  tar,  and  soluble  materials 
depends  upon  many  factors,  such  as  the  composition  of  the 
coal,  the  temperature  employed,  and  the  rate  of  heating.  One 
ton  of  good  gas  coal  yields  approximately  10,000  cu.  ft.  of  gaa, 
1400  Ib.  of  coke,  120  Ib.  of  tar,  and  20  gal.  of  animoniacal  liquor. 


FIG.  116.   Diagram  of  a  plant  used  for  the  manufacture  of  coal  gas 
and  its  by-products 

Not  only  is  the  ammonia  obtained  in  the  manufacture  of 
the  gas  of  great  importance  (p.  205)  but  the  coal  tar  is  the 
source  of  many  very  useful  substances ;  for  example,  most 
dyes,  as  well  as  most  of  the  explosives  used  in  war,  are  pre- 
pared from  compounds  derived  from  coal  tar.  This  subject 
will  be  discussed  in  a  later  chapter. 

The  by-product  coke  oven.  It  will  be  observed  that  coke  is 
formed  in  the  process  used  in  the  manufacture  of  coal  gas. 
Coke  is  a  very  important  product  and  is  used  in  large  quanti- 
ties, especially  in.  the  reduction  of  metals,  such  as  iron,  from 
their  ores.  The  quantity  of  coke  obtained  in  the  manufacture 
of  coal  gas  has  never  been  sufficient  to  meet  the  demand.  The 
additional  coke  required  has  been  prepared  for  the  most  part 


308    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

by  coking  the  coal  in  ovens  called  beehive  ovens  because  of 
their  shape.  The  coking  of  coal  in  these  ovens  is  carried  out 
as  follows  :  The  oven  is  nearly  filled  with  coal  and  the  coal  is 
ignited.  After  the  fire  is  well  started  the  draft  is  shut  off, 
and  the  heat  formed  in  the  combustion  of  a  portion  of  the  coal 
is  sufficient  to  coke  the  remainder  of  the  coal.  In  this  process 
all  the  coal  tar,  coal  gas,  and  ammonia  escape  through  an 
opening  in  the  top  of  the  furnace  and  are  lost.  The  growing 
demand  for  ammonia,  as  well  as  for  the  products  obtained  from 
coal  tar,  has  led  to  the  construction  of  furnaces  or  ovens  for 
the  coking  of  coal  which  make  it  possible  to  save  the  coal  tar 
and  ammonia  formed  in  the  process.  Such  ovens  are  known  as 
by-product  coke  ovens,  this  term  being  chosen  because  the  am- 
monia and  coal  tar  formed  in  the  process  of  coking  the  coal 
are  by-products,  the  coke  being  the  main  product.  These  ovens 
are  very  much  more  complex  than  the  beehive  ovens,  but  the 
demand  for  ammonia  and  coal  tar  is  rapidly  becoming  so  great 
as  to  cause  the  gradual  introduction  of  the  by-product  ovens, 
although  most  of  our  coke  is  still  made  in  the  beehive  ovens. 
Fig.  117  represents  a  large  by-product  coking  plant.  The 
ovens  A  are  placed  side  by  side  in  an  upright  position.  The 
coal  is  carried  up  through  the  shaft  B  and  introduced  into 
the  ovens  from  above.  The  coal  tar  collects  in  the  large  main 
C.  The  more  volatile  portions  which  escape  condensation  in  C 
are  condensed  in  appropriate  pipes.  A  portion  of  the  coal  gas 
generated  in  the  process  is  used  as  fuel  for  coking  the  coal. 

Water  gas.  Water  gas  is  essentially  a  mixture  of  carbon 
monoxide  and  hydrogen.  It  is  manufactured  by  passing 
superheated  steam  over  very  hot  anthracite  coal  or  coke, 
the  chief  reactions  being  expressed  in  the  following 
equations : 

C  +  H2O  — >•  CO  +  H2  -  26,990  cal. 
C02  +  C  — *•  2  CO  -  37,230  cal. 

The  industrial  process  is  intermittent.  The  fuel  is  burned 
with  a  forced  draft  in  a  suitable  furnace  until  it  is  very  hot. 


FUELS;  FLAMES;  ELECTRIC  FURNACES      309 

The  air  is  then  shut  off  and  the  steam  turned  on  until  the 
temperature  falls  to  about  1000°.  The  process  is  then  reversed. 
The  fall  in  temperature  is  rapid,  partly  owing  to  radiation  and 
to  the  cooling  occasioned  by  the  steam,  but  largely  because  of 
the  endothermic  character  of  the  reactions  which  take  place. 
The  gas  so  formed  contains  all  the  nitrogen  which  was  in  the 
furnace  when  the  steam  was  admitted. 

Water  gas  burns  with  a  pale-blue  nonluminous  flame. 
It  is  very  poisonous  and  has  no  odor.    To  make  it  suitable 


FIG.  117.   A  modern  by-product  coke  oven' 

for  illumination  in  an  ordinary  burner,  as  well  as  to  give 
it  an  odor  and  so  make  it  safer,  it  must  be  enriched  with 
hydrocarbons  called  illuminants.  This  is  accomplished  by 
passing  the  gas  through  a  furnace  filled  with  hot  fire 
brick  upon  which  crude  petroleum  is  sprayed.  The  petro- 
leum oils  are  decomposed  (cracked)  into  simpler  gaseous 
bodies,  the  most  important  of  which  are  methane,  acety- 
lene, and  ethylene.  Coal  gas  is  sometimes  enriched  in  a 
similar  way  by  adding  petroleum  to  the  coal  in  the  retorts. 
Producer  gas.  Producer  gas  is  used  in  connection  with 
many  metallurgical  furnace  operations  and  also  as  a  fuel 
for  gas  engines.  It  is  made  by  burning  coal  under  such 


310    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


conditions  that  the  product  of  combustion  is  largely  carbon 
monoxide  (Fig.  118).  Very  often  a  little  steam  is  admitted 
with  the  air,  and  this  on  passing  through  the  hot  bed  of 
coals  is  reduced  as  in  the  preparation  of  water  gas.  Made  in 
this  way,  producer  gas  is  composed  mainly  of  carbon  mon- 
oxide, hydrogen,  and  nitrogen.  It  can  be  made  from  coal 

of  a  poor  quality,  even 
from  lignite,  and  as  gas 
engines  run  well  with 
this  gas,  it  furnishes  the 
most  economical  method 
for  utilizing  low-grade 
coal  for  power. 

Natural  gas.  In  many 
regions  of  the  United 
States,  as  well  as  in 
other  countries,  natural 
gas  is  obtained  from 
wells  drilled  into  a 
stratum  holding  the 
gas.  While  it  is  va- 
riable in  composition, 
it  consists  largely  of 
methane,  many  samples  containing  as  much  as  95  per  cent 
of  this  compound.  It  burns  with  a  flame  of  moderate  lumi- 
nosity, but  works  well  with  a  gas  mantle.  It  has  a  high 
heat  of  combustion,  as  shown  in  the  following  equation : 

CH4  +  2  O2  — *  CO2  +  2  H2O  +  213,500  cal. 

It  is  an  ideal  fuel  and  is  often  conducted  through  pipes 
for  hundreds  of  miles  from  the  gas  fields  to  cities. 

Comparative  composition  of  gases.    The  following  figures  are 
the  results  of  analyses  of  average  samples,  but  since  each  kind 


FIG.  118.     Diagram   of   the  method 
making  producer  gas 


FUELS;  FLAMES;  ELECTRIC  FURNACES      311 

of  gas  varies  considerably  in  composition,  the  values  are  to  be 
taken  as  approximate  only.  The  nitrogen  and  traces  of  oxygen 
are  derived  from  the  air. 


COMPOSITION  OF  GASES  EXPRESSED  IN  PERCENTAGE  BY  VOLUME 


OHIO 

ENRICHED 

CONSTITUENT  GASES 

NATURAL 
GAS 

COAL 

GAS 

GAS 

WATER 
GAS 

GAS 

H2  

41.3 

52.88 

37.96 

10.90 

CH,, 

89  5 

43.6 

2.16 

7  09 

9  3 

2  01 

0.3 

3.9 

9.40 

0.60 

CO  

0.4 

6.4 

36.80 

32.25 

20.10 

CO2      

0.3 

2.0 

3.47 

4.73 

8.50 

N2   

0.2 

1.2 

4.69 

3.96 

59.90 

o 

03 

0  60 

Other  hydrocarbons 

1.5 

1.80 

Relation  of  the  two  gases  to  the  flame.  The  gas  issuing 
from  the  burner  is  said  to  undergo 
combustion,  while  that  one  which 
constitutes  the  atmosphere  about 
the  flame  is  said  to  support  com- 
bustion. These  terms  are  entirely 
conventional,  since  the  relation  of 
the  two  gases  may  be  reversed  with- 
out greatly  altering  the  appearance 
of  the  flame. 

Fig.  119  illustrates  a  convenient 
apparatus  for  demonstrating  this  fact. 
A  wide  lamp  chimney  (A~)  is  covered 
with  a  piece  of  asbestos  board  (£) 
which  has  a  hole  in  the  center  about  as 
large  as  a  dime.  A  straight  tube  (C) 
about  1  cm.  wide  and  also  a  smaller 


FIG.  119.  Apparatus  for 
showing  the  relation  of 
two  gases  to  the  flame 


tube  (D)  connected  with  the  gas  supply     produced  by  their  union 


312    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

pass  through  a  cork  at  the  bottom.  If  the  hole  in  B  is  closed 
(by  a  piece  of  asbestos  board)  while  gas  is  admitted  through  D, 
the  excess  gas  escapes  downward  through  C,  where  it  may  be 
lighted.  If  the  hole  in  B  is  now  opened,  the  flame  will  ascend 
to  the  top  of  the  tube  C.  This  flame  is  produced  by  air  drawn 
up  through  C,  burning  in  an  atmosphere  of  coal  gas.  Finally, 
the  excess  of  coal  gas  may  be  ignited  at  B,  where  it  will  burn 
in  air,  the  two  flames  being  very  similar  in  their  appearance. 


FIG.  120.   Materials  used  in  making  gas  mantles,  and  stages  in  the 
process  of  manufacture 

Gas  mantles.  In  using  the  fuel  gases  as  illuminants  the 
gas  is  usually  mixed  with  air  before  burning.  In  this  way 
the  gas  burns  with  a  hot  but  nearly  nonluminous  flame. 
The  light  is  obtained  by  suspending  about  this  flame  a 
gauze  mantle  of  suitable  material.  The  best  mantles  are 
composed  of  a  mixture  of  99  per  cent  of  thorium  oxide 
with  1  per  cent  of  cerium  oxide. 

The  thorium  and  cerium  compounds  used  in  gas  mantles 
are  obtained  from  monazite  sand  (Fig.  120)  found  prin- 
cipally in  North  Carolina  and  Brazil.  The  process  of 


FUELS;  FLAMES;  ELECTRIC  FURNACES      313 

making  a  gas  mantle  consists  in  knitting  a  tubular  fabric, 
which  is  then  dipped  into  a  solution  of  the  nitrates  of 
thorium  and  cerium.  After  being  dried  the  fabric  is 
heated,  in  which  process  the  yarn  is  burned,  while  the 
nitrates  of  thorium  and  cerium  are  converted  into  oxides 
which  are  left  in  the  form  of  the  original  fabric.  The 
resulting  mantle  is  very  delicate  and  is  strengthened  for 
shipping  by  dipping  it  into  a  solution  of  an  appropriate 
substance  and  drying. 

Products  of  the  combustion  of  ordinary  fuels.  Ordinary 
fuels,  such  as  oil,  wood,  coal,  and  fuel  gases,  are  largely 
made  up  of  carbon  and  hydrogen  or  their  compounds. 
The  chief  products  of  the  combustion  of  such  fuels  are 
carbon  dioxide  and  water.  Associated  with  these  are  small 
amounts  of  other  products,  such  as  carbon  monoxide  and 
sulfur  dioxide,  the  later  being  formed  from  traces  of  sulfur 
compounds  in  the  fuels. 

Rooms  are  not  infrequently  heated  by  gas  or  oil  stoves, 
with  no  provisions  for  removing  the  products  of  combus- 
tion. Likewise,  natural  gas  is  often  burned  in  stoves  or 
grates  with  the  damper  closed  so  as  to  leave  no  opening 
into  the  chimney.  Such  practices  are  greatly  to  be  con- 
demned, since  the  air  in  the  rooms  heated  in  this  way 
soon  becomes  so  contaminated  with  the  various  products 
of  combustion  as  to  render  it  unfit  for  respiration.  The 
large  amount  of  water  vapor  formed  in  rooms  so  heated 
condenses  on  the  windows  in  cold  weather,  causing  the 
glass  to  sweat. 

Conditions  necessary  for  flames.  When  one  of  the  sub- 
stances undergoing  combustion  remains  solid  at  the  tem- 
perature occasioned  by  the  combustion,  light  may  be  given 
off,  but  there  is  no  flame.  Thus,  iron  wire  burning  in 
oxygen  throws  off  a  shower  of  sparks,  but  no  flame  is  seen. 


314    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


When,  however,  both  of  the  substances  involved  are  gases 
or  vapors  at  the  temperature  reached  in  the  combustion, 
the  act  of  union  is  accompanied  by  a  flame. 

Flames  from  burning  liquids  or  solids.  Many  sub- 
stances which  are  liquids  or  solids  at  ordinary  tempera- 
tures burn  with  a  flame  because  the  heat  of  combustion 
slowly  vaporizes  them,  and  the  flame  is  due  to  the 
union  of  this  vapor  with  the  oxygen  of  the  air.  This 
may  be  shown  in  the  case 
of  a  candle  flame  by  hold- 
ing one  end  of  a  slender 
glass  tube  in  the  base  of  the 
flame  (Fig.  121).  The  un- 
burned  vapor  in  the  inner 
part  of  the  flame  is  thus 
conducted  away,  and  may 
be  ignited  at  the  upper  end 
of  the  tube. 

Bunsen    burners.      In    the 


contains  combustible  gases  in  similar  burners  used  in  gas 
ranges  and  for  illumination 

with  the  aid  of  mantles,  the  gas  is  mixed  with  a  certain 
percentage  of  air  before  it  is  burned.  This  is  accom- 
plished by  having  an  opening  (mixer)  in  the  base  of 
the  burner  into  which  the  air  is  drawn  by  the  flow  of 
the  gas.  If  the  inixer  is  adjusted  so  that  the  proper 
amount  of  air  is  admitted,  the  flame  is  colorless.  Such 
a  flame  possesses  an  advantage  in  that  it  is  very  hot  and 
no  carbon  is  deposited  from  it. 

Structure  of  a  flame.  The  structure  of  a  flame  can  be 
studied  to  the  best  advantage  when  the  combustible  gas 
issues  from  a  round  tube  into  an  atmosphere  of  the  gas 


FUELS;  FLAMES;  ELECTRIC  FURNACES      315 


supporting  combustion  (usually  the  air),  as  is  the  case 
with  an  ordinary  Bunsen  burner  (Fig.  122).  Under  these 
conditions  the  flame  is  conical  in  outline. 
Simple  flames.  When  the  chemical 
action  taking  place  in  the  combustion 
is  the  mere  union  of  two  gases,  as  is 
true  in  the  union  of  hydrogen  or  carbon 
monoxide  with  oxygen  or  of  hydrogen 
with  chlorine,  the  structure  of  the  flame 
is  very  simple.  It  consists  of  two  super- 
imposed cones  of  different  altitudes.  The 
inner  one  may  be  shown  to  be  merely 
unchanged  cold  gas,  and  is 
therefore  not  a  real  part  of 
the  flame.  A  match  head 
suspended  in  this  region 
(Fig.  122)  before  lighting 
the  gas  is  not  ignited  by 
the  flame  around  it. 

Complex  flames.  In  the  burning  of  hydro- 
carbons, as  well  as  of  many  other  gases,  the 
flame  is  more  complex,  and  as  many  as  four 
distinct  cones  may  be  seen  (Fig.  123).  The 
innermost  one  (^4)  is  really  not  a  part  of  the 
flame,  being  formed  of  gas  not  yet  brought 
to  the  point  of  combustion.  If  a  Bunsen 
burner  is  employed,  with  the  ring  at  the 
base  turned  to  admit  plenty  of  air,  the  sec- 
ond cone  (Z?)  is  sharply  defined  and  is  bluish  FIG.  123.  The 

green  in  color.    If  the  burner  tube  is  wide     conesof  acom- 
,     .       ,     ,  ,  plex  flame 
or  too  much  air  is  admitted,  the  rate  or  com- 
bustion in  this  cone  may  exceed  the  rate  of  flow  of  the 
gas,  in  which  case  cone  A  will  disappear  and  the  flame 


FIG.    122.   A  sim- 
ple Bunsen  flame 


316    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


will  travel  down  the  tube  and  burn  at  the  base,  or  strike 
back.  As  the  air  is  shut  off  it  will  be  seen  that  a  lumi- 
nous spot  appears  at  the  apex  of  cone  B,  which  gradually 
takes  the  form  of  a  cone  (<7)  quite  covering  the  inner 
one  and  brightly  luminous  over  all  its  surface.  Finally, 
if  some  object  is  held  so  as  to  intercept  the  light  from 
this  region,  it  will  be  seen  that  there  is  a  fourth  cone  (Z>), 
which  is  only  faintly  luminous. 

In  the  candle  flame  (Fig.  1 24)  there 
are,  broadly  speaking,  three  cones: 
(1)  the  inner  cone  (.4),  composed  of 
combustible  vapors;  (2)  an  intermedi- 
ate cone  (5),  in  which  these  vapors  are 
decomposed  by  the  heat  and  carbon 
is  set  free,  which  renders  the  flame 
luminous ;  and  (3)  an  almost  invisi- 
ble narrow  outer  cone,  or  film  (£),  in 
which  the  carbon  and  hydrogen  are 
burned  to  water  and  carbon  dioxide. 
Luminosity  of  flames.  As  the  cold 
gas  in  the  inner  cone  moves  toward 
the  hottest  region  of  the  flame,  its 
temperature  rapidly  rises,  and  at  definite  temperatures, 
which  depend  upon  the  nature  of  the  gas,  decomposition 
takes  place.  These  decompositions  cause  sharp  changes 
in  the  density  of  the  gas,  and  this  in  turn  makes  the  gas 
visible,  just  as  air,  heated  by  a  hot  pavement  in  summer, 
is  visible  in  wavering  lines.  In  this  decomposition,  prod- 
ucts may  be  formed  that  are  solids  (such  as  carbon),  and 
these  become  incandescent  at  the  temperature  reached, 
making  the  flame  still  more  luminous.  Each  chemical 
change  really  produces  a  distinct  cone  in  a  steady  flame. 
The  luminosity  of  a  flame  thus  depends  upon  many  factors. 


FIG.  124.   The  cones  of 
a  candle  flame 


FUELS;  FLAMES;  ELECTRIC  FURNACES      317 


The  following  equations  show  some  of  the  successive 
reactions  that  may  take  place  in  the  combustion  of  three 
different  gases : 


CH4 

H2S 

2H3As 


-C+2H2 

-S+H2 


S+H20  —  >-  S02+  H,0 


-2As+3H2 — >-2As+3H2O 


•As203+3H20 

If  we  have  a  flame  from  any  one  of  these  burning  gases 
and  suddenly  place  a  cold  object  (a 
small  porcelain  dish)  in  the  flame,  the 
free  element  is  chilled  below  its  kin- 
dling temperature  and  is  deposited  as 
soot  (carbon,  sulfur,  arsenic)  upon 
the  dish. 


-1540' 


1550° 


--1560' 


1570- 


1450°- 


-\    -1540° 


-    -520' 


The  temperature  of  flames.  The  actual 
temperature  which  can  be  realized  in  an 
ordinary  flame  obviously  depends  upon 
many  conditions,  such  as  the  composi- 
tion of  the  gas,  its  pressure,  tempera- 
ture, and  rate  of  flow,  and  the  method  of 
supplying  the  air.  Even  in  an  ordinary 
Bunsen  flame  burning  under  favorable 
conditions  it  is  very  difficult  to  deter- 
mine the  maximum  temperature  attained. 
The  actual  region  of  great  heat  is  lim- 
ited, as  the  burning  zones  are  very  thin. 
The  temperature  in  different  parts  of 
the  flame  varies  greatly,  and  any  object 
placed  in  the  flame  for  determining  its 
temperature  cuts  across  many  different 
regions  and  is  unequally  heated.  Evi- 
dently  the  temperature  is  much  higher 
than  that  recorded  by  a  body  in  the 

flame,  since  the  specific  heat  of  solids  is  so  much  greater  than 
that  of  gases.  Under  exceptional  conditions  it  has  been  found 


FIG.  125.   An  estimate 

of  the  temperature  in 

various     parts     of     a 

Bunsen  flame 


318    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

possible  to  melt  a  very  fine  platinum  wire  in  a  good  Bunsen 
flame,  so  that  a  temperature  of  1755°  is  surely  reached.  The 
accompanying  diagram  (Fig.  125)  gives  a  rough  estimate  of  the 
probable  temperature  in  various  parts  of  a  good  nonluminous 
Bunsen  flame. 

Reducing  and  oxidizing  flames.    Since  the  region  just 
below  the  luminous  cone   is  very  hot  and  contains  the 


FIG.  126.    Diagram  of  a  smoke-consuming  furnace 

reducing  gases  hydrogen  and  carbon  monoxide,  a  sub- 
stance such  as  a  metallic  oxide  placed  in  this  region  will 
undergo  reduction,  provided  it  can  be  reduced  by  such 
hot  gases.  A  body  heated  in  this  way  is  said  to  be  heated 
in  the  reducing  flame.  At  the  apex  of  the  flame  there  are 
no  reducing  gases,  but  it  is  very  hot  and  air  is  abundant ; 
consequently  a  substance  which  is  rather  readily  oxidized 
will  undergo  oxidation  if  heated  in  this  region.  This  part 
of  the  flame  is  called  the  oxidizing  flame. 


FUELS;  FLAMES;  ELECTRIC  FURNACES      319 


Smoke  prevention.  Since  the  products  of  combustion  of 
fuels  are  carbon  dioxide  and  water  vapor  and  these  are 
invisible  compounds,  it  is  evident  that  if  the  combustion 
is  complete  no  smoke 
will  be  formed.  As  a 
rule  the  combustion 
is  imperfect ;  gaseous 
compounds  containing 
carbon  are  first  formed, 
and  when  these  are 
imperfectly  burned  a 
part  of  their  carbon 
is  set  free  in  a  finely 
divided  state  consti- 
tuting smoke.  Smoke 
may  therefore  be  pre- 
vented by  securing  the 
complete  combustion 
of  the  fuel,  the  neces- 
sary conditions  being 
as  follows:  (1)  a  suffi- 
cient supply  of  air; 
(2)  thorough  mixing 
of  the  ah'  with  the  combustible  gases  produced  from  the 
fuel ;  and  (3)  a  temperature  high  enough  to  maintain 
active  combustion. 

Smoke  prevention  is  a  problem,  of  great  economic  impor- 
tance, especially  in  the  large  cities.  Thus,  for  example,  it  has 
been  estimated  that  the  smoke  in  the  city  of  Pittsburgh  costs 
the  people  of  the  city  $10,000,000  yearly,  or  about  $20  for 
each  inhabitant ;  and  this  does  not  take  into  account  the  pos- 
sible effect  of  smoke  upon  health.  Because  of  these  facts  many 
cities  are  now  taking  steps  to  abate  the  smoke  nuisance.  That 


FIG.  127.   A  bomb  calorimeter 


320    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

the  conditions  necessary  for  preventing  smoke  may  be  met,  it 
is  essential  that  the  coal  be  introduced  into  the  furnace  uni- 
formly, so  that  the  volatile  matter  expelled  upon  heating  may 
be  more  readily  mixed  with  air.  This  is  done  efficiently  by 
having  a  chain  grate,  as  is  shown  in  A  (Fig.  126).  The  coal  is 
fed  on  this  at  B,  and  as  the  chain  slowly  moves  forward,  the 
coal  gradually  enters  the  furnace,  and  by  the  time  it  reaches  the 
back  part  of  the  furnace  (C)  it  is  completely  burned,  the  ashes 
falling  out  at  D.  The  volatile  matter  expelled  is  thoroughly 
mixed  with  hot  air  led  in  through  the  back  of  the  grate  E,  E. 
The  large  space  under  the  boiler  drum  gives  opportunity  for 
complete  combustion  of  the  products  under  the  chimney.  The 
water  in  the  drum  circulates  through  the  tubes  as  shown  by 
the  arrows  and  thus  is  heated  to  a  high  temperature. 

Calorific  value  of  fuels.  The  various  materials  used  as 
fuels  differ  much  in  the  heat  which  they  give  out  when 
burned.  While  many  other  factors  are  concerned  in  the 
value  of  a  fuel,  the  chief  one  is  its  heat  of  combustion. 
The  heat  evolved  by  the  combustion  of  one  gram  of  a 
fuel  is  called  its  calorific  value.  In  large  contracts  the 
price  paid  for  a  fuel  is  generally  based  on  its  calorific 
value,  as  well  as  upon  its  adaptability  to  the  use  to  which 
it  is  to  be  put.  The  following  table  will  give  some 
average  values  for  a  few  common  fuels : 

CALORIFIC  VALUE  OF  FUELS 

Wood  (air-dried) about  3800-4000  cal. 

Lignite  (brown),  8%  ash,  12%  moisture  ....  about  5400  cal. 
Bituminous  coal  (Pennsylvania),  35%  volatile 

matter,  6%  ash about  8300  cal. 

Bituminous  coal  (Pocahontas),  18%  volatile 

matter,  6%  ash about  8700  cal. 

Anthracite  coal  (Connellsville),  12%  ash  ....  about  7300  cal. 

Coke,  10%  ash about  7300  cal. 


FUELS;  FLAMES;  ELECTRIC  FURNACES      321 


In  determining  the  calorific  value  of  a  fuel  an  instrument 
known  as  a  bomb  calorimeter  is  used.  This  is  a  strong  steel 
flask  lined  with  platinum  or  porcelain  and  provided  with  a 
tight-fitting  screw  cap  (Fig.  127).  In  determining  the  heat  of 
combustion  a  weighed  sample  of  the  substance  is  placed  on  the 
capsule  A,  oxygen  is  admitted  through  the  tube  B  until  the 
pressure  in  the  bomb  is  about  20  atmospheres,  and  the  bomb 
is  then  closed  and  placed  in  an  open  calorimeter.  The  charge 
is  ignited  by  passing  an  electric  current  through  the  fine  iron 
fuse-wire  C  stretched  above  the  charge.  The  wire  is  melted, 
and  the  red-hot  drop  of  burning  metal  falls  upon  the  charge, 
igniting  it.  The  heat  given  off  dur- 
ing combustion  is  measured  by  the 
rise  in  temperature  of  the  water 
surrounding  the  bomb,  which  is 
stirred  by  the  stirrer  D.  A  pre- 
liminary experiment  must  be  made 
upon  a  weighed  charge  of  a  sub- 
stance whose  heat  of  combination 
is  known  (such  as  cane  sugar), 
to  determine  the  heat  absorbed 
by  the  bomb,  together  with  that  due  to  the  melting  and  com- 
bustion of  the  fuse-wire,  and  also  the  loss  by  radiation. 

The  electric  furnace.  In  recent  years  electric  furnaces  have 
come  into  wide  use  in  operations  requiring  a  very  high  tem- 
perature. Temperatures  as  high  as  3500°  can  be  easily  reached, 
whereas  the  hottest  oxyhydrogen  flame  is  not  much  above 
2000°.  These  furnaces  are  constructed  on  one  of  two  general 
principles. 

1.  Arc  furnaces.  In  the  one  type  the  source  of  heat  is  an 
electric  arc  formed  between  carbon  electrodes  separated  a  little 
from  each  other,  as  shown  in  Fig.  128.  The  substance  to  be 
heated  is  placed  in  a  vessel,  usually  a  graphite  crucible,  just 
below  the  arc.  The  electrodes  and  crucible  are  surrounded  by 
materials  which  fuse  with  great  difficulty,  such  as  magnesium 
oxide,  the  walls  of  the  furnace  being  so  shaped  as  to  reflect 
the  heat  downwards  upon  the  contents  of  the  crucible. 


FIG.  128.   An  arc  furnace 


322    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

2.  Resistance  furnaces.  In  the  other  type  of  furnace  the  heat 
is  generated  by  the  resistance  offered  to  the  current  in  its 
passage  through  the  furnace.  In  its  simplest  form  it  may  be 
represented  by  Fig.  129.  The  furnace  is  merely  a  rectangular 
box  built  up  of  loose  bricks.  The  electrodes  (E,  E),  each  consist- 
ing of  a  bundle  of  carbon  rods,  are  introduced  through  the  sides 
of  the  furnace.  The  materials  to  be  heated  (C)  are  filled  into  the 

furnace  up  to  the  elec- 
trodes, and  a  layer  of 
broken  coke  is  arranged 
so  as  to  extend  from  one 
electrode  to  the  other. 
More  of  the  charge  is 
FIG.  129.  A  resistance  furnace  then  placed  on  top  of 

the    coke.      In    passing 

through  the  broken  coke  the  electrical  current  encounters  great 
resistance.  This  generates  much  heat,  and  the  charge  surround- 
ing the  coke  is  brought  to  a  very  high  temperature.  The 
advantage  of  this  type  of  furnace  is  that  the  temperature  can 
be  regulated  to  any  desired  intensity. 


EXERCISES 

1.  Why  does  charcoal  usually  burn  with  no  flame  ?    How  do  you 
account  for  the  flame  sometimes  observed  when  it  burns  ? 

2.  How  do  you  account  for  the  fact  that  a  candle  burns  with 
a  flame? 

3.  What  two  properties  must  the  mantle  used  in  the  Welsbach 


4.  (a)  In  what  respects  does  the  use  of  the  Welsbach  mantle 
resemble  that  of  lime  in  the  calcium  light?    (b)  If  the  mantle  were 
made  of  carbon,  would  it  serve  the  same  purpose  ? 

5.  Would  anthracite  coal  be   suitable  for  the  manufacture  of 


6.  How  could  you  prove  the  formation  of  carbon  dioxide  and 
water  in  the  combustion  of  illuminating  gases  ? 

7.  Suggest  a  probable  way  in  which  natural  gas  has  been  formed. 


FUELS;  FLAMES;  ELECTRIC  FURNACES     323 

8.  Coal  frequently  contains  a  sulfide  of  iron,    (a)  What  two 
sulfur  compounds  are  likely  to  be  formed  when  gas  is  made  from 
such  coal?    (6)  Suggest  some  suitable  method  for  the  removal  of 
these  compounds. 

9.  Why  does  the  use  of  the  bellows  on  the  blacksmith's  forge 
cause  a  more  intense  heat? 

10.  What  name  is  applied  to  reactions  such,  as  those  which  take 
place  in  the  manufacture  of  water  gas? 

11.  Water  gas  was  used  as  a  fuel  in  a  stove  which  consumed 
5  cu.  ft.  of  gas  per  hour.    Calculate  the  volume  of  oxygen  required 
to  burn  the  gas,  the  fuel  and  oxygen  being  measured  under  the 
same  conditions  of  temperature  and  pressure. 

12.  Suppose  that  natural  gas  is  pure  methane  and  that  a  stove 
burns  20 1.  of  the  fuel  per  hour,  measured  at  20°  and  740  mm. 
(a)  What  volume  of  oxygen  measured  at  20°  and  740  mm.  would  be 
required  for  the  combustion?    (//)  If  a  room  were  heated  by  the 
stove  and  no  arrangement  was  made  for  carrying  off  the  products 
of  combustion,  what  volume  of  carbon  dioxide  (20°  and  740  mm.) 
would   be  added   to  the   air  in  the  room?     (e)   What  weight  of 
moisture  would  be  formed? 


CHAPTER  XXVI 
CARBOHYDRATES ;  ALCOHOLS  ;  COAL-TAR  COMPOUNDS 

Carbohydrates.  The  term  carbohydrate  is  applied  to  a 
class  of  compounds  which  includes  the  sugars,  starch,  and 
allied  substances.  These  compounds  contain  carbon,  hydro- 
gen, and  oxygen,  the  last  two  elements  usually  being  pres- 
ent in  the  proportion  in  which  they  combine  to  form  water. 
The  most  important  carbohydrates  are  the  following: 

TABLE  OF  CARBOHYDRATES 

Sucrose  (ordinary  sugar) C12H22On 

Lactose  (milk  sugar) C12H22On  •  H2O 

Maltose C12H22On  •  H2O 

Dextrose  (grape  sugar) C6H12O6 

Levulose C6H12O6 

Cellulose (C6H10O5)X 

Starch (C6H1006)X 

The  molecular  formulas  of  cellulose  and  starch  are  un- 
known, but  are  multiples  of  the  simple  formula  C6H1QO5; 
accordingly  they  are  often  written  (C6H10O5)X.  In  the 
discussion  of  the  compounds  they  will  be  represented  by 
the  simple  formula  C6H10O5. 

It  will  be  noted  that  some  of  the  compounds  named  in 
the  table  have  identical  formulas.  Such  compounds  are 
said  to  be  isemeric.  The  difference  in  the  properties  of 
isomeric  compounds  is  due  to  the  fact  that  the  atoms  are 
arranged  differently  in  the  molecule. 


CARBOHYDRATES  325 

Sucrose  (sugar)  (C12H22011).  This  substance,  commonly 
called  sugar,  occurs  in  many  plants,  especially  in  the  sugar 
cane  and  sugar  beet,  each  of  which  at  present  furnishes 
about  50  per  cent  of  the  total  production.  The  sugar  cane 
grows  only  in  warm  climates  (Cuba  and  the  Hawaiian 
Islands  are  the  greatest  producers),  while  the  sugar  beet 
thrives  in  cooler  climates,  such  as  prevail  in  Ohio  and 
Michigan  in  the  United  States,  and  in  Germany.  The 
beets  contain  as  high  as  16  per  cent  of  sucrose. 

The  manufacture  of  sugar.  The  juice  from  the  cane  or  beet 
contains  the  sugar  iu  solution  along  with  many  impurities. 
These  impurities  arc  removed  by  appropriate  methods  and  the 
resulting  solution  is  then  evaporated  until  the  sugar  crystal- 
lizes. The  evaporation  is  conducted  in  closed  vessels  from 
which  the  air  is  partially  exhausted  (vacuum  pans).  In  this 
way  the  boiling  point  of  the  solution  is  lowered  and  the  char- 
ring of  the  sugar  is  prevented.  It  is  not  practicable  to  remove 
all  the  sugar  from  the  solution.  Ordinary  molasses  is  the  solu- 
tion which  remains  after  a  part  of  the  sugar  has  been  crys- 
tallized out  from  the  purified  juice  of  the  sugar  cane.  The 
sweetness  of  maple  sugar  is  due  to  sucrose,  other  products 
present  in  the  maple  sap  imparting  the  distinctive  flavor. 
About  40,000,000,000  Ib.  of  sugar  is  produced  annually.  The 
annual  consumption  of  sugar  in  the  United  States  is  over 
9,000,000,000  Ib.,  or  approximately  90  Ib.  for  each  person. 

Chemical  conduct  of  sugar.  When  a  solution  of  cane 
sugar  is  heated  to  about  70°  with  hydrochloric  acid,  two 
isomeric  sugars,  dextrose  and  levulose,  are  formed  in  accord- 
ance with  the  following  equation  : 

<W>,,  +  H20  — »  C0H120,  +  C,HU0, 

In  this  process  the  sugar  is  said  to  be  inverted,  and  the 
mixture  of  dextrose  and  levulose  is  termed  invert  sugar. 


326    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

When  heated  to  160°  sucrose  melts ;  if  the  temperature 
is  increased  to  about  215°  a  partial  decomposition  takes 
place  and  a  brown  substance  known  as  caramel  is  formed. 
This  is  used  extensively  as  a  coloring  matter  and  in  making 
confectionery. 

Lactose  (milk  sugar)  (C12H22011  •  H20).  This  compound 
is  present  in  the  milk  of  all  mammals.  The  average 
composition  of  cow's  milk  is  as  follows: 

Water 87.17% 

Casein  (nitrogenous  matter) 3.56% 

Butter  fat 3.64% 

Lactose 4.88% 

Mineral  matter 0.75% 

When  rennin  (a  substance  obtained  from  the  stomach  of 
calves)  is  added  to  milk  the  casein  separates.  This  is  the 
part  of  the  milk  used  in  the  manufacture  of  cheese.  The 
liquid  remaining  after  the  separation  of  the  casein  is  known 
as  whey.  This  contains  the  milk  sugar,  which  crystallizes 
on  evaporation ;  it  resembles  sucrose  in  appearance,  but  is 
not  so  sweet  nor  so  soluble.  The  souring  of  milk  is  due  to 
the  fact  that  the  milk  sugar  contained  in  it  changes  into 
lactic  acid,  a  liquid  having  the  formula  C3H6O3: 

CBHM0U  +  H,0->4C.H.O, 

This  change  is  brought  about  by  a  certain  microorganism 
which  enters  from  the  air,  and  the  process  is  known  as 
lactic  fermentation. 

Dextrose  (grape  sugar,  glucose)  (C6H1206).  This  sugar  is 
present  in  honey  and  in  many  fruits.  It  is  usually  asso- 
ciated with  levulose,  and  is  often  called  grape  sugar  because 
of  its  presence  in  grape  juice.  It  can  be  obtained  along 
with  levulose  by  heating  sucrose  with  hydrochloric  acid,  as 
explained  above.  Commercially  it  is  prepared  in  enormous 


CARBOHYDRATES  32T 

quantities  by  heating  starch  with  hydrochloric  acid.  The 
starch  is  first  changed  into  a  sweet-tasting  solid  known 
as  dextrin,  and  this,  on  further  action,  is  converted  into 
dextrose :  c  „  Q  ,  H  Q  _^  c  „  Q 


10 


When  the  change  is  complete  the  hydrochloric  acid  is 
neutralized  by  sodium  carbonate.  Over  50,000,000  bushels 
of  corn  are  used  each  year  in  the  United  States  in  the 
production  of  dextrose  and  allied  products. 

Pure  dextrose  is  a  white  crystalline  solid  resembling 
sucrose  in  its  properties,  but  not  so  sweet.  Most  of  the 
dextrose  used  is  in  the  form  known  commercially  as  glucose, 
or  corn  sirup.  This  is  a  thick  sirupy  liquid  and  consists 
of  an  aqueous  solution  of  dextrin,  dextrose,  and  maltose. 
Large  quantities  of  glucose  are  used  in  the  preparation  of 
jellies,  jams,  sirups,  candy,  and  other  sweets.  A  federal 
ruling  requires  that  when  glucose  is  present  in  such  foods 
as  jellies  and  jams,  the  label  on  the  container  must  state 
the  percentage  of  glucose  present. 

Starch  (C6H1005).  This  substance  is  always  present  in 
seeds  and  tubers  and  is  by  far  the  most  abundant  carbo- 
hydrate found  in  nature.  In  the  United  States  it  is  obtained 
chiefly  from  corn,  about  60  per  cent  of  which  is  starch.  In 
Europe  the  potato  serves  as  the  principal  source. 

The  manufacture  of  starch.  In  manufacturing  starch  from 
corn,  the  corn  is  first  soaked  in  water  containing  a  little  sul- 
furous  acid,  to  soften  the  grain.  It  is  then  ground  coarsely  so 
as  not  to  crush  the  germ.  When  the  resulting  mass  is  mixed 
with  water  the  germ  floats,  being  very  light  because  of  the  oil 
which  it  contains.  In  this  way  the  germ  is  separated  from  the 
rest  of  the  seed,  and  from  it  corn  oil  is  prepared.  The  remain- 
ing material,  consisting  of  the  starch,  the  nitrogenous  con- 
stituent (gluten),  and  the  bran,  or  outside  coating  of  the  grain, 
is  then  ground  finely,  mixed  with  water,  and  passed  through 


328    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

cloth  sieves,  which  remove  the  bran.  The  water  containing  the 
starch  and  gluten  in  suspension  is  then  allowed  to  run  slowly 
down  long,  shallow  troughs,  the  rate  of  flow  being  so  regulated 
that  the  heavier  starch  sinks  to  the  bottom  of  the  trough  while 
the  lighter  gluten  is  washed  away.  The  starch  is  then  removed 
from  the  troughs  as  shown  in  Fig.  130  and  subsequently  dried. 
Large  quantities  of  starch  are  used  in  making  glucose  and 
other  foods,  for  finishing  .cloth,  and  for  laundry  purposes. 


FIG.  130.   Removing  the  starch  from  the  settling  troughs  in  a 
starch  factory 

Characteristics  of  starch.  Starch  consists  of  minute  gran- 
ules, which  differ  somewhat  in  appearance  according  to 
the  source  of  the  starch,  so  that  it  is  often  possible  from  a 
microscopic  examination  to  determine  from  what  plant  any 
given  sample  of  starch  was  obtained  (Figs.  131  and  132). 
When  heated  with  water  the  granules  burst  and  the 
starch  partially  dissolves.  This  is  the  reason  why  starchy 
foods  are  made  more  digestible  by  cooking. 

Cellulose  (C6H1005).  Cellulose  forms  the  basis  of  all 
woody  fibers.  Cotton  and  linen  are  nearly  pure  cellulose. 
It  is  insoluble  in  water,  alcohol,  or  dilute  acids,  but  will 


CARBOHYDRATES 


329 


FIG.  131.   Wheat-starch  gran- 
ules magnified  200  diameters 


dissolve  in  a  solution  prepared  by  dissolving  copper  oxide 
in  ammonium  hydroxide.  Concentrated  hydrochloric  acid 
changes  it  hi  to  dextrose.  Concentrated  nitric  acid  forms  a 
mixture  of  compounds  which  is  known  as  nitrocellulose  or 
guncotton.  These  are  very  in- 
flammable and  under  certain 
conditions  are  highly  explosive. 
They  have  many  commercial 
uses.  Photographic  films  are 
made  from  them,  as  well  as 
from  a  noninflammable  deriva- 
tive of  cellulose  known  as  acetyl 
cellulose.  Collodion  is  a  solution  of 
certain  nitrocelluloses  in  a  mix- 
ture of  alcohol  and  ether.  Celluloid  is  a  mixture  of  nitrocellu- 
lose and  camphor.  These  two  when  mixed  together  form  a 
plastic  mass  whicli  can  be  molded  into  any  desired  shape  and 
which  is  used  for  making  such  objects  as  combs  and  brush 
handles.  Celluloid  is  very  inflammable  and  care  should  be 
exercised  in  the  use  of  celluloid  articles. 

Mercerized  cotton  and  artificial  silk. 
When  cotton  cloth  is  treated  with  a  con- 
centrated solution  of  sodium  hydroxide, 
the  cellulose  shrinks  and  becomes  tougher 
in  character.  If  the  cloth  is  placed  in 
stretchers  to  prevent  the  shrinkage,  it 
assumes  an  appearance  somewhat  resem- 
bling silk  and  is  known  as  mercerized 
cotton.  Another  fabric  prepared  in  large 
quantities  from  cellulose  resembles  silk 

very  closely  and  is  known  as  artificial  silk.  The  fiber  of  this 
fabric  is  prepared  by  forcing  concentrated  solutions  of  cellu- 
lose or  its  derivatives  through  minute  tubes  and  coagulating 
the  cellulose  as  it  emerges  in  the  form  of  fine  threads. 


FIG.  132.     Cornstarch 

granules  magnified  200 

diameters 


830    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Characteristics  of  various  textile  fibers.  Of  the  different 
fibers  used  in  making  the  yarns  from  which  the  common 
fabrics  are  prepared,  the  vegetable  fibers,  cotton  and  linen, 
are  essentially  cellulose,  while  the  animal  fibers,  wool  and 
silk,  are  composed  of  nitrogenous  substances.  Although 
these  fibers  resemble  each  other  when  viewed  with  the 
naked  eye,  their  appearance  is  distinctive  when  examined 
with  the  microscope.  The  characteristic  appearance  of 
these  fibers  is  shown  in  Fig.  133.  It  is  also  possible  to 
distinguish  between  the  fibers  by  the  action  of  chemical 
reagents.  For  example,  a  hot  solution  of  sodium  hydroxide 


Silk  fiber  Cotton  fiber  Wool  fiber 

FIG.  133.   Three  important  textile  fibers 

(5  per  cent  to  10  per  cent)  has  but  little  action  upon  cotton, 
while  it  will  readily  dissolve  wool  and  slowly  dissolve  silk. 
Paper.  Paper  consists  mainly  of  cellulose.  The  finer 
grades  are  made  from  linen  and  cotton  rags  and  the  cheaper 
grades  from  wood. 

Manufacture  of  paper.  In  making  paper  the  raw  material  is 
cut  into  pieces  and  treated  with  suitable  reagents  (calcium  acid 
sulfite  is  used  in  case  of  wood),  to  remove  all  objectionable 
matter,  leaving  the  cellulose,  which  is  then  bleached  with  chlo- 
rine. The  paper  pulp  so  obtained  is  suspended  in  water  and 
run  onto  wire  screens.  It  then  passes  between  large  iron 
cylinders,  some  of  which  are  heated  with  steam.  In  this  way 
the  pulp  is  pressed  and  dried  and  delivered  in  the  form  of 


ALCOHOLS 


331 


paper.  In  the  process  different  materials  are  often  added  to 
the  pulp.  These  vary  with  the  nature  of  the  paper  desired; 
thus,  finely  ground  clay  or  calcium  sulfate  is  added  to  give  body 
to  the  paper.  In  making  paper  intended  for  writing  or  print- 
ing, a  compound  prepared  by  heating  resin  and  sodium 
hydroxide  is  added,  together  with  aluminium  sulfate.  This 
makes  a  finished  surface  and  prevents  the  ink  from  spreading. 


FIG.  134.   The  interior  of  a  paper  mill 

Fig.  134  shows  the  interior  of  a  paper  mill.  The  pulp  flows 
from  the  container  A  onto  the  screens  beyond  and  then  be- 
tween the  rollers  until  it  is  pressed  and  dried  and  so  converted 
into  the  finished  paper  B. 


ALCOHOLS 

The  alcohols  may  be  regarded  as  derived  from  the  hydro- 
carbons by  substituting  for  one  or  more  hydrogen  atoms 
a  corresponding  number  of  hydroxyl  groups.  A  great 
many  alcohols  are  known,  and,  like  the  hydrocarbons,  they 


332    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

may  be  arranged  in  series.  The  relation  between  the  first 
three  members  of  the  methane  series  of  hydrocarbons  and 
the  corresponding  alcohols  is  shown  in  the  following  table : 

CH4  (methane) CH8OH  (methyl  alcohol) 

C2H6  (ethane) C2H5OH  (ethyl  alcohol) 

C3H8  (propane) C3H7OH  (propyl  alcohol) 

The  terms  methyl,  ethyl,  and  propyl,  used  in  designating  the 
different  alcohols,  are  names  applied  to  the  univalent  radicals 
CH8,  C2H5,  and  C3H7  respectively.  It  will  be  noted  that  the 
names  of  these  radicals  are  derived  from  the  names  of  the  cor- 
responding hydrocarbons  by  changing  the  ending  -one  to  -yl. 

Methyl  alcohol  (wood  alcohol)  (CH3OH).  This  compound 
is  formed  when  wood  is  heated  in  the  absence  of  air 
(p.  121),  and  on  this  account  it  is  called  wood  alcohol.  It 
is  a  colorless  liquid  which  boils  at  64.7°  and  burns  with 
an  almost  colorless  flame.  It  is  a  good  solvent  for  organic 
substances  and  is  used  extensively  in  the  manufacture  of. 
varnishes.  It  is  poisonous.  It  also  has  a  specific  action  on 
the  optic  nerve,  and  many  persons  have  become  blind  from 
drinking  the  liquid  or  from  repeatedly  inhaling  its  vapor. 

When  a  mixture  of  the  vapor  of  methyl  alcohol  and  air  is 
passed  over  hot  copper,  the  alcohol  is  oxidized,  forming  a 
gaseous  compound  known  as  formaldehyde : 

2  CH3OH  +  02 >•  2  CH20  +  2  H2O 

This  gas  is  now  prepared  in  large  quantities  and  used  as  a 
disinfectant.  A  40  per  cent  aqueous  solution  of  it  is  sold 
under  the  name  of 


Ethyl  alcohol  (grain  alcohol,  alcohol)  (C2H5OH).  This 
compound  is  the  one  ordinarily  known  as  alcohol.  It 
resembles  methyl  alcohol  in  its  general  properties. 


ALCOHOLS 


333 


1.  Preparation.    It  is  prepared  by  the  action  of  ordinary 
baker's  yeast  upon  different  sugars,  such  as  dextrose: 

C  H  O  — >•  2  C  HKOH  +  2  CO, 

6       12      6  26  2 

This  process  in  which  a  sugar  is 
changed  into  alcohol  and  carbon 
dioxide  by  the  action  of  yeast  is 
known  as  alcoholic  fermentation. 
The  yeast  is  a  low  form  of  plant 
life  (Fig.  135)  and  thrives  in  ap- 
propriate sugar  solutions.  During 

its  growth  a  number  of  changes 
FIG.  135.    Some  cells  of  the     take      ]ace   which   regult   ^    COQ_ 
yeast  plant  . 

verting  the  sugar  into  alcohol. 

Experimental  preparation  of  alcohol.  The  formation  of  alcohol 
and  carbon  dioxide  from  dextrose  may  be  shown  as  follows : 
A  10  per  cent  solution  of  the  sugar  in  water  is  poured  into 
flask  A  (Fig.  136)  and  a  little  baker's  yeast  is  added.  The 
bottle  B,  containing  limewater,  is  connected  as  shown  in  the 
figure.  The  tube  C 
is  filled  with  pieces 
of  sodium  hydroxide 
to  prevent  carbon 
dioxide  from  entering 
from  the  air.  The 
temperature  is  main- 
tained at  about  30°. 
Action  soon  begins, 
as  is  indicated  by 
the  bubbles  of  carbon 
dioxide,  and  continues 
until  the  sugar  is  all  FIG.  136.  Laboratory  preparation  of  alcohol 
fermented.  That  the 

escaping  gas  is  carbon  dioxide  is  shown  by  the  precipitate 
formed  in  B.   The  alcohol  formed  is  separated  by  distillation. 


334    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Commercial  preparation  of  alcohol.  Alcohol  is  prepared  com- 
mercially from  starch  obtained  from  corn  or  potatoes.  The 
starch  is  first  converted  into  a  sugar  known  as  maltose,  by 
the  action  of  malt,  a  substance  prepared  by  moistening  barley 
with  water,  allowing  it  to  sprout,  and  then  drying  it.  Maltose, 
like  dextrose,  breaks  down  into  alcohol  and  carbon  dioxide  in 
the  presence  of  yeast.  The  resulting  alcohol  is  separated  by 
fractional  distillation. 

2.  Properties.  Ethyl  alcohol  is  a  colorless  liquid  with  a 
pleasant  odor  and  is  an  excellent  solvent  for  many  organic 
substances.  It  boils  at  78.3°.  It  is  sometimes  used  as 
a  fuel,  since  its  flame  is  very  hot  and  does  not  deposit 
carbon,  as  the  flame  from  oil  does.  When  taken  into 
the  system  in  small  quantities  it  causes  intoxication ;  in 
larger  quantities  it  acts  as  a  poison.  The  ordinary  alco- 
hol of  the  druggist  contains  about  95  per  cent  alcohol 
and  5  per  cent  water.  A  solution  containing  99  per  cent 
or  more  of  alcohol  is  called  absolute  alcohol.  When  alcohol 
is  heated  with  sulfuric  acid  a  low-boiling  inflammable 
liquid  known  as  ether  13  formed : 

2C2H6OH— >(C2H5)20  +  H20 

This  is  largely  used  as  an  anesthetic  in  surgical  operations 
and  as  a  solvent  in  chemical  industries. 

Denatured  alcohol.  While  ordinary  alcohol  (95  per  cent) 
is  not  difficult  to  prepare,  its  selling  price  is  greatly  in- 
creased because  of  the  high  tax  imposed  by  the  government. 
By  an  act  of  Congress  in  1906  the  tax  was  removed  from 
denatured  alcohol;  that  is,  alcohol  mixed  with  some  sub- 
stance which  renders  it  unfit  for  use  as  a  beverage  but 
does  not  impair  its  use  for  manufacturing  purposes.  The 
substances  ordinarily  used  for  this  purpose  are  methyl 
alcohol  and  benzine. 


ALCOHOLS  335 

Alcoholic  liquors.  All  alcoholic  liquors  are  made  by  alcoholic 
fermentation.  'Wine  is  made  by  the  fermentation  of  the  dex- 
trose in  grape  juice  and  contains  from  5  to  15  per  cent  by 
volume  of  alcohol.  Beer  is  made  from  maltose  formed  by  the 
action  of  malt  upon  starch  obtained  from  various  grains,. chiefly 
barley.  It  contains  from  3  to  5  per  cent  by  volume  of  alcohol. 
Whisky  contains  about  50  per  cent  by  volume  of  alcohol  and 
is  made  from  starch  by  a  process  very  similar  to  that  described 
under  the  commercial  preparation  of  alcohol.  Almost  any  sac- 
charine liquid,  such  as  cider  and  the  juices  of  fruits  in  general, 
undergoes  alcoholic  fermentation  when  exposed  to  air. 

Alcoholic  liquors,  as  well  as  pure  alcohol,  are  taxed  by  the 
government.  Previous  to  the  passage  of  the  prohibition  amend- 
ment over  $200,000,000  was  collected  annually  from  this  source. 

Chemical  changes  in  bread-making.  The  average  com- 
position of  wheat  flour  is  as  follows: 

Water 13.8% 

Protein  (nitrogenous  matter) 7.9% 

Fats 1.4% 

Starch 76.4% 

Mineral  matter 0.5% 

In  making  bread,  flour  is  mixed  with  water,  yeast,  and 
a  little  sugar,  and  the  resulting  dough  is  set  aside  in  a 
warm  place  for  a  few  hours.  The  yeast  first  causes  the 
sugar  to  undergo  alcoholic  fermentation.  The  carbon  diox- 
ide formed  escapes  through  the  dough,  making  it  light  and 
porous.  The  yeast  plant  thrives  best  at  about  30°  ;  hence 
the  necessity  for  keeping  the  dough  in  a  warm  place.  In 
baking  bread  the  beat  expels  the  alcohol  and  also  expands 
the  bubbles  of  carbon  dioxide  caught  in  the  dough,  causing 
it  to  become  porous  and  making  the  bread  light. 

Preservatives.  We  have  observed  that  the  changes  taking 
place  in  the  souring  of  milk  and  the  changing  of  sugar  into 
alcohol  are  caused  by  organisms  the  cells  of  which  are  present 


336    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

in  the  air.  Many  other  similar  changes,  such  as  putrefaction, 
are  due  to  the  same  causes.  All  these  change's  may  be  pre- 
vented in  one  of  the  following  ways  : 

1.  By  keeping  the  substance  at  such  a  low  temperature 
that  the   organism  ca.using   the  change   cannot  thrive   (cold 
storage). 

2.  The  substance  may  be  heated  so  as  to  destroy  all  organ- 
isms present  and  then  sealed  air-tight  in  a  suitable  container. 
This  is  the  method  used  in  canning  vegetables  and  in  preserv- 
ing such  substances  as  grape  juice  and  condensed  milk. 

3.  Some  substance  may  be  added  which  in  small  amounts 
will  destroy  the  organisms  causing  the  change  or  will  prevent 
their  growth.    Such  a  substance  is  known  as  a  preservative. 

Whether  or  not  preservatives  should  be  permitted  in  foods 
is  a  much-debated  question.  Some  people  maintain  that  any 
substance  which  is  powerful  enough  to  prevent  the  growth  of 
the  organisms  must  have  an  injurious  action  upon  digestion. 
The  federal  government  at  present  allows  the  use  of  sodium 
benzoate  (a  white  solid  made  from  a  hydrocarbon  present  in 
coal  tar)  in  such  foods  as  jellies,  jams,  and  catchup,  which 
are  not  consumed  immediately  upon  the  opening  of  the  con- 
tainer. If  this  preservative  is  used,  however,  the  labels  on 
the  containers  must  state  the  amount  present. 

Some  derivatives  of  coal  tar.  In  discussing  the  manu- 
facture of  coal  gas  (p.  306)  it  was  stated  that  from  the 
coal  tar  formed  in  the  process  there  is  obtained  a  large 
number  of  important  compounds.  These  are  often  spoken 
of  collectively  as  the  coal-tar  compounds.  It  is  possible  here 
to  mention  only  a  few  of  these. 

(1)  Benzene  (C6H6)  and  (2)  toluene  (C7Hg)  are  highly  inflam- 
mable colorless  liquids ;  (3)  naphthalene  (C^Hg)  and  (4)  an- 
thracene (CMH10)  are  white,  solid  hydrocarbons  which  are  used 
in  the  preparation  of  the  two  dves  indigo  and  alizarin.  These 
dyes  were  formerly  obtained  from  vegetable  sources,  but  are 
now  manufactured  at  low  cost.  Ordinary  moth  balls  are  nearly 


COAL-TAR  COMPOUNDS  337 

pure  naphthalene.  (5)  Phenol,  or  carbolic  acid  (C6H5OH),  is  a 
white  crystalline  solid,  very  caustic  and  poisonous.  (6)  Cresol 
(C.H7OH)  is  a  good  disinfectant  and  is  the  basis  of  most  of 
the  disinfectants  now  on  the  market. 

Each  of  the  above  compounds  serves  as  the  source  material 
from  which  many  other  useful  compounds  are  prepared.  Thus, 
benzene  when  treated  with  nitric  acid  gives  nitrobenzene 
(C6H5N02),  and  this  on  reduction  yields  aniline  (C6H5NH2). 
Aniline  is  a  nearly  colorless  liquid,  and  from  it  are  prepared  a 
large  number  of  dyes  of  all  shades  and  colors,  known  as  the 
aniline  dyes.  Toluene  when  oxidized  forms  benzole  acid,  the 
sodium  salt  of  which  (sodium  benzoate)  is  used  as  a  food  pre- 
servative. When  phenol  is  heated  with  formaldehyde  there  are 
obtained  products  known  commercially  as  bakelite  and  condens- 
ite.  These  are  useful  materials  for  making  buttons,  umbrella 
handles,  pipestems,  and  insulators  in  electrical  apparatus. 

Coal-tar  compounds  in  foods.  Much  discussion  has  arisen 
in  regard  to  the  use  of  coal-tar  compounds  in  foods.  It  is 
evident  that  no  substance  which  acts  injuriously  upon  the 
human  system  should  be  used  in  our  foods ;  neither 
should  the  la\v  permit  the  use  of  any  substance  which  is 
•used  for  purposes  of  deception.  The  federal  government 
has  selected  seven  aniline  dyes  of  different  colors,  the  use 
of  which  is  permitted  in  such  foods  as  candies  and  butter. 
The  use  of  sodium  benzoate  as  a  preservative  is  allowed 
under  certain  restrictions.  Saccharine,  a  white  solid  pre- 
pared from  toluene  and  500  times  as  sweet  as  sugar,  was 
formerly  permitted  in  foods,  but  in  1912  the  government 
forbade  its  further  use.  Vanillin,  identical  with  the  com- 
pound prepared  from  vanilla  beans,  and  coumarin,  which 
has  an  odor  similar  to  vanillin,  are  both  used  in  artificial 
vanilla  extracts,  but  when  they  are  so  used  the  label  on 
the  container  must  state  the  fact.  It  is  well  to  keep  in 
mind  that  these  substances  have  no  nutritive  value. 


338    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

EXERCISES 

1.  What  is  the  meaning  of  the  term  carbohydrate"}  (See  dictionary.) 

2.  Can  you  tell  the  difference  between  pure  sugar  obtained  from 
sugar  cane  and  that  obtained  from  the  sugar  beet  ? 

3.  It   is    often  said  that   milk    sours   readily    during  thunder 
showers.    What  would  you  say  as  to  the  truth  of  this  statement? 

4.  Why  do  we  use  corn  rather  than  dextrose  in  making  alcohol  ? 

5.  How  could  you  tell  the  difference  between  methyl  alcohol 
and  ethyl  alcohol? 

6.  Yeast  is  often  added  in  preparing  household  beverages  such 
as  root  beer.    Why  is  it  added?    What  substance  will  be  present  in 
the  beverage  so  prepared  ? 

7.  Why  is  sugar  (or  molasses)  added  in  making  bread  ? 

8.  Alcohol  and  gasoline  boil  at  about  the  same  temperature  and 
both  are  combustible.    Why  not  use  alcohol  as  a  fuel  in  place  of 
gasoline  ? 

9.  Can  you  suggest  a  method  for  obtaining  ethyl  alcohol  from 
wood? 

10.  What  weight  of  starch  is  necessary  in  making  100  kg.  of  pure 
dextrose  ? 

11.  1  kg.  of  sucrose  would  yield  what  weight  of  invert  sugar? 

12.  What  weight  of  dextrose  is  necessary  for  the  preparation  of 
10  kg.  of  the  ordinary  alcohol  of  the  druggist,  on  the  supposition 
that  95  per  cent  of  the  sugar  undergoes  fermentation  ? 

13.  What  weight  of  benzene  is  necessary  for  the  preparation  of 
100  kg.  of  aniline  ? 

14.  What  volume  of  carbon  dioxide  is  evolved  in  the  fermenta- 
tion of  100  g.  of  dextrose  ? 

15.  What  weight  of  methyl  alcohol  would  be  required  for  the 
preparation  of  50  kg.  of  formaldehyde  ? 

16.  Calculate  the  weight  of  lactic  acid  formed  in  the  souring  of 
10  kg.  of  milk. 


CHAPTER  XXVII 
ORGANIC  ACIDS ;   FATS  AND  OILS 

Organic  acids.  A  great  number  of  acids  are  known 
which  are  composed  of  carbon,  oxygen,  and  hydrogen, 
and  as  a  group  these  are  called  organic  acids.  Like  the 
hydrocarbons,  they  can  be  arranged  in  series,  one  of  the 
most  important  of  which  is  known  as  the  fatty-acid  series. 
A  few  of  the  most  important  members  of  this  series  are 
given  in  the  following  table.  They  are  all  monobasic  —  a 
fact  indicated  in  the  formula  by  separating  the  replaceable 
hydrogen  atom  from  the  rest  of  the  molecule. 

SOME  FATTY  ACIDS 

H-CHO2 formic  acid,  a  liquid  boiling  at  100° 

H  •  C2H3O2 acetic  acid,  a  liquid  boiling  at  118° 

H  •  C4H7O2 butyric  acid,  a  liquid  boiling  at  163° 

H  •  C16H31O2 palmitic  acid,  a  solid  melting  at  62° 

H  •  C18H85O2 stearic  acid,  a  solid  melting  at  69° 

H-CnH2n_!O2 general  formula 

Of  these  acetic  acid  deserves  special  mention. 

Acetic  acid  (H  •  C2H302).  This  is  the  acid  which  gives 
the  sour  taste  to  vinegar.  It  is  prepared  commercially  by 
the  distillation  of  wood  (p.  120).  It  is  a  colorless  liquid 
and  has  a  strong,  pungent  odor.  When  anhydrous  it 
crystallizes  as  a  white  solid  which  melts  at  17°  and  closely 
resembles  ice  in  appearance ;  hence  the  name  glacial  acetic 
acid.  Many  of  the  salts  of  acetic  acid  are  well-known 
compounds.  Thus,  lead  acetate  (Pb(C2H8O2)2 .  3  H.O)  is 
the  white  solid  known  as  sugar  of  lead. 


340    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Vinegar.  As  is  well  known,  when  cider  is  exposed  to  the  air 
it  is  gradually  transformed  into  vinegar.  Two  changes  are 
involved  in  the  process  :  (1)  the  sugar  in  the  cider  first  under- 
goes alcoholic  fermentation,  forming  hard  cider,  which  contains 
from  4  to  8  per  cent  of  alcohol ;  (2)  the  alcohol  is  then  oxidized 
to  acetic  acid,  the  necessary  oxygen  coming  from  the  air.  This 
oxidation  is  brought  about  through 
the  action  of  the  microorganism 
known  as  Mycoderma  aceti.  This 
organism  is  present  in  the  so-called 
mother  of  vinegar.  The  oxidation 
of  alcohol  into  acetic  acid  through 
the  agency  of  the  Mycoderma  aceti 
is  known  as  acetic  fermentation,  and 
may  be  represented  as  follows  : 

C2H6OH  +  02 >•  H  •  C2H302  +  H20 

The  manufacture  of  vinegar.  The 
old  method  of  making  vinegar  con- 
sisted simply  in  storing  cider  in 
barrels  until  the  fermentation  was 
complete.  In  the  modern  method  a 
large  cask,  known  as  a  generator,  is 
used  (Fig.  137).  This  is  filled  loosely 
with  beech  shavings.  Vinegar  is  first 
sprayed  into  the  top  of  the  cask  in 
order  to  introduce  the  Mycoderma 
aceti.  The  organism  attaches  itself 

to  the  wood  shavings,  which  are  used  because  they  present  a 
large  surface.  Next  a  dilute  solution  of  alcohol  (hard  cider, 
in  the  case  of  cider  vinegar)  is  sprayed  into  the  top  of  the 
cask  while  air  is  admitted  at  the  bottom  A,  A.  In  this  way 
the  alcohol  and  oxygen  are  brought  into  intimate  contact,  and 
the  oxidation  takes  place  rapidly  as  the  liquid  trickles  down 
over  the  shavings.  The  resulting  vinegar  is  drawn  off  at  the 
bottom  (B)  of  the  cask.  Instead  of  starting  with  cider,  one 
may  use  almost  any  substance  which  contains  starch  or  sugar, 


FIG.  137.    A  generator  for 
the  manufacture  of  vinegar 


ORGANIC  ACIDS;  FATS  AND  OILS  341 

these  compounds  first  being  changed  into  alcohol,  as  explained 
in  the  manufacture  of  alcohol.  In  this  way  are  prepared 
malt  vinegar  from  starch  and  sugar  vinegar  from  sugar  resi- 
dues. The  cheapest  vinegar  is  made  from  pure  dilute  alcohol, 
and  is  known  as  distilled  vinegar.  It  is  colorless  and  leaves  no 
residue  upon  evaporation. 

A  federal  law  requires  that  all  vinegar  shall  contain  not 
less  than  4  per  cent  acetic  acid.  In  addition  to  the  acid,  vine- 
gar prepared  from  fruits  and  grains  contains  certain  solids 
derived  from  the  source  materials.  It  is  by  studying  the  char- 
acter of  these  solids  left  upon  evaporating  a  sample  of  vinegar 
that  the  chemist  is  able  to  determine  the  source  of  the  vinegar. 

Acids  belonging  to  other  series.  In  addition  to  the  fatty 
acids,  the  following  deserve  special  mention: 

Tartaric  acid  (H2  •  C4H406).  This  is  a  white  solid  and 
occurs  in  many  fruits,  either  in  the  free  state  or  in  the 
form  of  its  salts.  The  acid  potassium  salt  KHC4H4O6 
occurs  in  the  juice  of  grapes.  When  the  juice  ferments 
in  the  manufacture  of  wine,  this  salt,  being  insoluble  in 
alcohol,  is  deposited  on  the  sides  of  the  cask,  in  which 
form  it  is  known  as  argol.  When  purified  it  forms  a  white 
solid,  which  is  sold  under  the  name  of  cream  of  tartar  and 
is  used  in  baking  powders.  The  acid  itself  is  often  used 
in  soft  drinks. 

Citric  acid  (H3  •  C6HB07).  This  acid  occurs  in  citrus 
fruits,  such  as  lemons  and  grape  fruit.  It  is  a  white  solid, 
soluble  in  water. 

Oleic  acid  (H  •  ClgH3302).  The  derivatives  of  this  acid 
constitute  the  principal  part  of  many  oils  and  liquid  fats. 
The  acid  itself  is  an  oily  liquid. 

Fats  and  oils.  The  hydrogen  of  acids  can  be  replaced 
not  only  by  metals  but  by  hydrocarbon  radicals  as  well. 
The  resulting  compounds  are  termed  esters.  The  main 
constituents  of  the  common  fats  and  oils,  such  as  butter, 


342    AN  ELEMENTAKY  STUDY  OF  CHEMISTRY 

lard,  and  olive  oil,  are  esters  of  oleic,  palmitic,  and  stearic 
acids  and  are  known  respectively  as  olein,  palmitin,  and 
stearin.  The  radical  present  in  these  esters  is  CgH5.  It 
is  trivalent  and  is  known  as  the  glyceryl  radical,  since 
it  is  present  in  glycerin  (C8H5(OH)3).  Since  the  glyceryl 
radical  is  trivalent,  and  since  oleic,  palmitic,  and  stearic 
acids  are  all  monobasic,  it  is  evident  that  three  molecules 
of  each  acid  must  enter  into  the  formation  of  each  mole- 
cule of  the  ester  derived  from  it.  The  relation  in  com- 
position between  these  acids  and  the  corresponding  esters 
is  shown  in  the  following  formulas : 

H  •  C18H83O2  (oleic  acid) 'C3H5(C18H33O2)3  (olein) 

H  •  C16H31O2  (palmitic  acid)  .  .  .  C3H5(C16H31O2)3  (palmitin) 
H  •  C18H3502  (stearic  acid)  ....  C3H5(C18H35O2)3  (stearin) 

Olein  is  a  liquid  and  is  the  main  constituent  of  oils  such 
as  olive  oil.  Palmitin  and  stearin  are  white  solids.  Beef 
suet  is  principally  stearin. 

Butter  fat  and  oleomargarine.  While  butter  fat,  like 
other  fats,  consists  principally  of  olein,  palmitin,  and  stearin, 
its  characteristic  flavor  is  due  to  the  presence  of  a  small 
amount  (about  8  per  cent)  of  the  fat  butyrin,  which  is  an 
ester  of  butyric  acid  and  has  the  formula  CgH5(C4H7O2)3. 
Oleomargarine  differs  from  butter  mainly  in  the  fact  that  a 
smaller  amount  of  butyrin  is  present.  It  is  made  from  the 
fats  obtained  from  cattle  and  hogs.  Sometimes  cottonseed 
oil  is  also  added.  These  fats  are  churned  with  milk  or 
mixed  with  a  small  amount  of  butter,  in  order  to  furnish 
sufficient  butyrin  to  give  the  butter  flavor. 

In  appearance  oleomargarine  differs  from  most  butter  in 
being  nearly  colorless.  While  it  is  a  common  practice  to  color 
butter  artificially,  the  federal  law  permits  the  coloring  of  oleo- 
margarine only  upon  the  payment  of  a  tax  of  10  cents  for 


ORGANIC  ACIDS;  FATS  AND  OILS  343 

each  pound  colored.  Many  of  the  states,  however,  have  laws 
forbidding  the  sale  of  oleomargarine  that  is  artificially  colored, 
even  though  the  federal  tax  has  been  paid. 

Changing  oils  into  solid  fats.  It  will  be  noted  that  stearin 
differs  from  olein  in  composition  in  that  it  contains  six  more 
atoms  of  hydrogen  in  each  molecule.  Now  if  hydrogen  is 
brought  in  contact  with  olein  under  proper  conditions  and 
in  the  presence  of  a  suitable  catalytic  agent  (finely  divided 
nickel  is  used),  the  olein  takes  up  the  additional  hydrogen 
and  is  changed  into  the  solid  stearin.  It  is  possible  in  this 
way  to  change  the  oils  into  solid  fats.  Certain  commercial 
fats  used  in  cooking  are  made  by  this  process  from  the  com- 
paratively inexpensive  cottonseed  oil. 

The  proteins.  The  term  protein  is  applied  to  a  large  class 
of  complex  nitrogenous  compounds  which  are  everywhere 
abundant  in  animal  and  vegetable  organisms  and  which  con- 
stitute the  principal  part  of  the  tissues  of  the  living  cell. 
The  casein  of  milk,  gluten  of  flour,  and  albumin  of  egg 
will  serve  as  examples  of  protein  matter.  The  proteins  all 
contain  nitrogen,  carbon,  hydrogen,  and  oxygen,  and  some 
contain  sulfur  and  phosphorus  in  addition. 

Foods.  While  the  compounds  present  in  our  foods  are 
very  numerous  and  often  exceedingly  complex,  yet  they 
may  all  be  included  in  a  few  general  classes.  It  is  cus- 
tomary to  regard  the  edible  portion  of  our  foods  as  com- 
posed of  proteins,  fats,  carbohydrates,  mineral  matter, 
and  water.  Since  the  mineral  matter  is  left  as  a  residue 
when  the  food  is  burned,  it  is  listed  as  ash  in  reporting 
the  analyses  of  foods. 

In  a  general  way  it  may  be  stated  that  the  proteid 
matter  in  our  food  serves  to  replace  the  worn-out  tissues 
of  our  bodies,  as  well  as  to  'supply  material  for  growth. 
The  carbohydrates  and  fats  are  more  or  less  interchange- 
able, since  both  are  oxidized  in  the  body  and  serve  as  a 


344    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


source  of  heat  and  muscular  energy.  The  mineral  matter 
supplies  the  material  for  building  up  the  solid  tissues  of 
the  body  and  has,  in  addition,  other  more  complex  func- 
tions. The  proteid  matter  may  fulfill  the  same  function 
as  the  fats  and  carbohydrates  if  the  latter  are  lacking  in 
our  foods.  Since  the  various  constituents  of  our  foods 
serve  different  purposes,  it  is  evident  that  a  proper  mixture 
of  these  is  essential  to  health. 

The-  composition  of  the  edible  portion  of  a  few  typical 
foods  is  given  in  the  following  table,  taken  from  Sherman's 
"  Chemistry  of  Food  and  Nutrition." 


WATER 
(Per  cent) 

PROTEIN 
(Per  cent) 

FAT 

(Per  cent) 

CARBOHY- 
DRATES 
(Per  cent) 

ASH 

(Per  cent) 

Beef  free  from  visible  fat 
Ham,  smoked,  lean    .     . 
Salmon 

73.8 
53.6 
64.6 
73.7 
87.0 
11.0 
7.3 
12.3 
11.9 
35.3 
12.6 
75.4 
78.3 
94.3 
84.6 

22.1 
20.2 
21.2 
14.8 
3.3 
1.0 
16.1 
8.0 
13.3 
9.2 
22.5 
3.1 
2.2 
0.9 
0.4 

2.9 
20.8 
12.8 
10.5 
4.0 
85.0 
7.2 
0.3 
1.5 
1.3 
1.8 
1.1 
0.1 
0.4 
0.5 

5.0 

67.5 
79.0 
72.7 
53.1 
59.6 
19.7 
18.4 
3.9 
14.2 

1.2 
5.5 
1.4 
1.0 
0.7 
3.0 
1.9 
0.4 
0.6 
1.1 
3.5 
0.7 
1.0 
0.5 
0.3 

Eggs    
Milk    
Butter      

Oatmeal  
Rice    
Wheat  flour      .     . 
Bread,  white     .     .     .     . 
Beans,  dried      .... 
Corn,  green  

Potatoes  

Tomatoes      
Apples     

The  cycle  of  carbon  in  nature.  By  means  of  the  energy 
supplied  by  sunlight  the  carbon  dioxide  absorbed  from 
the  air  by  plants  reacts  with  water  and  small  amounts 
of  other  substances  derived  from  the  soil  to  form  com- 
plex compounds  of  carbon  which  constitute  the  essential 
part  of  the  plant  tissue.  This  reaction  is  attended  by 


ORGANIC  ACIDS;  FATS  AND  OILS  345 

the  evolution  of  oxygen,  which  is  restored  to  the  air.  The 
compounds  resulting  from  these  changes  are  much  richer 
in  energy  than  are  the  substances  from  which  they  are 
formed,  the  source  of  this  energy  being  the  sunshine. 

If  the  plant  is  burned  or  decays  in  the  open  air,  the 
changes  which  took  place  in  the  formation  of  the  con- 
stituents of  the  plant  are  largely  reversed.  The  carbon 
and  hydrogen  combine  with  oxygen  taken  from  the  air 
to  form  carbon  dioxide  and  water,  while  the  energy 
absorbed  from  the  sun's  rays  is  liberated  in  the  form  of 
heat.  If,  on  the  other  hand,  the  plant  is  used  as  food, 
the  compounds  present  are  utilized  in  building  up  the 
tissues  of  the  body  and  as  a  source  of  energy.  In  either 
case  the  carbon  and  hydrogen  ultimately  combine  with 
inhaled  oxygen  to  form  carbon  dioxide  and  water,  which 
are  in  turn  exhaled.  The  energy  possessed  by  the  food 
substance  is  liberated  partly  in  the  form  of  heat,  which 
maintains  the  temperature  of  the  body,  and  partly  as 
muscular  energy.  The  carbon  dioxide  originally  absorbed 
from  the  air  by  the  plant  is  thus  restored,  and  the  cycle 
of  changes  begins  anew. 


1.  For  what  purpose  have  we  used  formic   acid?    Consult  the 
dictionary  for  the  derivation  and  significance  of  the  word  formic. 

2.  What  weight  of  dextrose  would  be  required  to  prepare  100  kg. 
of  vinegar  containing  the  legal  quantity  of  acetic  acid,  on  the  sup- 
position that  all  the  sugar  is  converted  into  acetic  acid  ? 

3.  Account  for  the  fact  that  potassium-acid  tartrate   separates 
when  grape  juice  ferments. 

4.  Aluminium  is  a  trivalent  metal.   Write  the  formula  for  alu- 
minium stearate  ;  aluminium  tartrate  ;  aluminium  citrate. 


CHAPTER  XXVIII 
THE  PHOSPHORUS  FAMILY 


NAME  OF  ELEMENT 

SYMBOL, 

ATOMIC 

WEIGHT 

DENSITY 

MELTING 

POINT 

Phosphorus 

p 

31  04 

1  8 

44° 

Arsenic  

As 

74.96 

5.73 

850° 

Antimony  
Bismuth 

Sb 
Bi 

120.2 
208  0 

6.52 

9  8 

630° 
271° 

The  family.  The  elements  constituting  this  family  be- 
long '  in  the  same  group  with  nitrogen  and  therefore 
resemble  it  in  a  general  way  in  the  type  of  compounds 
formed.  They  exhibit  the  gradation  of  physical  properties 
shown  in  the  above  table.  The  same  general  gradation  is 
also  found  in  their  chemical  characteristics,  phosphorus 
being  an  acid-forming  element,  while  bismuth  is  essen- 
tially a  metal.  The  other  two  elements  are  intermediate 
in  character. 

Compounds.  In  general  the  elements  of  the  family  form 
compounds  having  similar  composition,  as  is  shown  in  the 
following  table: 

PC18  PC16  P2O3  P2O6 

AsCl3  AsCl5  As2O3  As2O6 

SbCl3  SbCl5  Sb203  Sb205 


PH3 

AsH8 

SbH3 


BiCls 


BiCl5 


Bi2O3 


In  the  case  of  phosphorus,  arsenic,  and  antimony  the 
oxides  are  acid  anhydrides.    Salts  of  at  least  four  acids  of 
each  of  these  three  elements  are  known,  the  free  acid  in 
346 


THE  PHOSPHORUS  FAMILY  347 

some  instances  being  unstable.  The  relation  of  these  acids 
to  the  corresponding  anhydrides  may  be  illustrated  as 
follows,  phosphorus  being  taken  as  an  example : 

P2O3  +  3  H2O >-2  H3PO3  (phosphorous  acid) 

P2O6  +  3  H2O >-2  H8PO4  (phosphoric  acid) 

P2O5  +  2  H2O >-    H4P2O7  (pyrophosphoric  acid) 

P2O6  +     H2O >•  2  HPO8  (metaphosphoric  acid) 

PHOSPHORUS 

History.  The  element  phosphorus  was  discovered  by 
the  alchemist  Brand,  of  Hamburg,  in  1669,  while  search- 
ing for  the  philosophers'  stone.  Owing  to  its  peculiar 
properties  and  the  secrecy  which  was  maintained  about 
its  preparation,  it  remained  a  very  rare  and  costly  sub- 
stance until  the  demand  for  it  in  the  manufacture  of 
matches  brought  about  its  production  on  a  large  scale. 

Occurrence.  Owing  to  its  great  chemical  activity  phos- 
phorus never  occurs  free  in  nature.  In  the  form  of 
phosphates  it  is  very  abundant  and  widely  distributed. 
Phosphorite  is  the  chief  mineral  form  of  calcium  phosphate, 
while  apatite  consists  of  calcium  phosphate  together  with 
calcium  fluoride  or  calcium  chloride.  These  minerals  form 
very  large  deposits  and  are  extensively  mined  for  use  as 
fertilizers.  Calcium  phosphate  is  a  constituent  of  all  fertile 
soil,  having  been  supplied  to  the  soil  by  the  disintegration 
of  rocks  containing  it.  It  is  the  chief  mineral  constituent 
of  the  bones  of  animals,  and  bone  ash  is  therefore  nearly 
pure  calcium  phosphate. 

Preparation.  Phosphorus  is  now  manufactured  from 
bone  ash  or  a  pure  mineral  phosphate  by  heating  the 
phosphate  with  sand  and  carbon  in  an  electric  furnace. 
Sand  consists  largely  of  silica  (SiO2),  and  this  is  the 


348    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


anhydride  of  metasilicic  acid  (H2SiO3).  At  very  high 
temperatures  phosphoric  anhydride  (P2O5)  is  volatile, 
while  at  the  same  temperature  silica  is  not.  Consequently 
when  calcium  phosphate  is  heated  with  silica  the  follow- 
ing equilibrium  is  set  up : 

Ca8(P04)2  +  3  Si02^=t  3  CaSi03  +  P2O5 

The  phosphoric  anhydride  formed  in  the  reaction  is  then 
reduced  by  the  carbon,  as  follows: 

2P2O5+10C >-P4+10CO 

The  materials  are  fed  in  at 
A  (Fig.  138)  by  the  feed  screw 
B.  The  phosphorus  vapor  es- 
capes at  D  and  is  condensed 
under  water,  while  the  calcium 
silicate  is  tapped  off  as  a  liquid 
at  C.  The  phosphorus  obtained 
in  this  way  is  quite  impure 
and  is  purified  by  distillation 
or  by  melting  it  and  pressing 
it  through  cloth. 

Properties.  The  purified 
Pl-Pl«>™.  called^,  or 
yellow,  phosphorus,  is  a  nearly 
colorless,  translucent,  waxy  solid  which  melts  at  44°  and 
boils  at  287°.  It  can  therefore  be  cast  into  any  conven- 
ient form  under  warm  water,  and  is  usually  sold  on  the 
market  in  the  form  of  sticks  (Fig.  139).  It  can  be  cut 
with  a  knife,  but  this  must  always  be  done  under  water, 
since  phosphorus  is  extremely  inflammable,  and  the  fric- 
tion of  the  knife  blade  is  almost  sure  to  set  it  on  fire  if 
it  is  cut  in  the  air.  It  is  not  soluble  in  water,  but  is  freely 
soluble  in  some  other  liquids,  notably  in  carbon  disulfide. 


THE  PHOSPHORUS  FAMILY 


349 


Its  density  is  1.8.  One  gram-molecular  volume  (22.4  1.) 
of  phosphorus  vapor  weighs  about  128  g.,  which  is  approx- 
imately four  times  the  atomic  weight,  showing  that  the 
formula  of  the  molecule  is  P4. 

Chemical  conduct.  When  exposed  to  the  air,  phosphorus 
slowly  combines  with  oxygen  and  in  so  doing  gives  out  a 
pale  light,  or  phosphorescence,  which 
can  be  seen  only  in  a  dark  place. 
The  heat  of  the  room  may  raise  the 
temperature  of  phosphorus  to  the 
kindling  point,  when  it  burns  with 
a  sputtering  flame,  giving  off  dense 
fumes  of  oxide  of  phosphorus.  It 
burns  with  dazzling  brilliancy  hi 
oxygen  and  combines  directly  with 
many  other  elements.  On  account 
of  its  great  attraction  for  oxygen 
it  is  preserved  under  water. 

Phosphorus  is  very  poisonous, 
from  0.2  to  0.3  g.  being  a  fatal 
dose.  Ground  with  flour  and  grease 
or  similar  substances,  it  is  used  as 
a  poison  for  rats  and  other  vermin. 

Red  phosphorus.  On  standing,  white  phosphorus  gradu- 
ally undergoes  a  remarkable  change,  being  converted  into 
a  dark -red  powder  which  has  a  density  varying  from  2.1  to 
2.38  and  which  is  called  red  phosphorus.  It  no  longer  takes 
fire  easily,  nor  is  it  soluble  in  carbon  disulfide.  It  is  not 
poisonous  and,  in  fact,  is  an  entirely  different  substance. 
The  velocity  of  this  change  of  white  phosphorus  to  red 
phosphorus  increases  with  rise  in  temperature,  and  red 
phosphorus  is  therefore  prepared  by  heating  the  white  form 
a  little  below  the  boiling  point.  When  distilled  and  quickly 


FIG.  139.    Sticks  of  white 
phosphorus 


350    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

condensed,  the  red  form  changes  back  to  the  white.  This 
is  in  accordance  with  the  general  rule  that  when  a  sub- 
stance capable  of  existing  in  several  forms  is  condensed 
from  a  gas  or  crystallized  from  the  liquid  state,  the  more 
unstable  variety  forms  first,  and  this  then  passes  into  the 
more  stable  forms. 

Matches.  The  chief  use  of  phosphorus  is  in  the  manu- 
facture of  matches,  two  general  varieties  of  which  are  in 
common  use.  Ordinary  friction  matches  are  made  by  dip- 
ping the  match  sticks  first  into  some  inflammable  substance, 
such  as  melted  paraffin,  and  afterward  into  a  paste  consisting 
of  (1)  phosphorus  sesquisulfide  (P4S3),  (2)  some  oxidizing 
substance,  such  as  manganese  dioxide  or  potassium  chlo- 
rate, and  (3)  a  binding  material,  usually  some  kind  of 
glue.  The  phosphorus  sulfide  is  ignited  by  friction,  the 
combustion  being  sustained  by  the  oxidizing  agent  and 
communicated  to  the  wood  by  the  burning  paraffin.  In 
sulfur  matches  the  paraffin  is  replaced  by  sulfur. 

In  safety  matches  a  mixture  of  red  phosphorus,  an  oxi- 
dizing agent,  and  some  gritty  material,  such  as  powdered 
glass,  is  placed  on  the  side  of  the  box,  while  the  match  tip 
is  provided  with  an  oxidizing  agent  and  an  easily  oxidizable 
substance,  usually  antimony  sulfide.  The  match  cannot  be 
ignited  easily  by  friction  except  on  the  prepared  surface. 

Matches  were  formerly  made  from  white  phosphorus,  and 
the  workmen  frequently  suffered  from  dreadful  diseases  of  the 
bones  of  the  face.  On  this  account  the  manufacture  and  use 
of  matches  containing  white  phosphorus  was  prohibited  in 
European  countries.  The  compound  P4Sg,  which  is  easily  pre- 
pared from  white  phosphorus,  serves  just  as  well  and  does  not 
occasion  disease,  and  in  1913  the  government  of  the  United 
States  placed  a  prohibitive  tax  (two  cents  per  hundred  matches) 
on  the  white  phosphorus  match,  at  the  same  time  forbidding 
both  the  import  and  the  export  of  such  matches. 


THE  PHOSPHORUS  FAMILY 


351 


Hydrides  of  phosphorus  —  phosphine.  Phosphorus  forms 
several  compounds  with  hydrogen,  the  best  known  of  which 
is  phosphine  (PH3)  analogous  to  ammonia  (NHg). 

The  simplest  way  of  making  phosphine  is  to  treat  calcium 
phosphide  with  water: 

Ca,P,  +  6  H20  — >-  3  Ca(OH)2  +  2  PH3 

It  is  more  conveniently  made  by  boiling  white  phosphorus 
suspended  in  a  concentrated  solution  of  sodium  hydroxide, 
the  reaction  being  a  complicated  one : 

P4  +  3  NaOH  +  3  H2O >-  3  NaH2PO2  +  PH3 

Phosphine  is  a  gas  of  unpleasant  odor  and  is  exceed- 
ingly poisonous.  Like  ammonia,  it  forms  salts  with  the 
halogen  acids.  Thus,  we  have  phosphonium  chloride 
(PH4C1)  analogous  to  ammonium  chloride  (NH4C1).  The 
phosphonium  salts  are  of 
but  little  importance. 

Phosphine  can  be  con- 
veniently made  in  the  ap- 
paratus shown  in  Fig.  140. 
A  concentrated  solution  of 
sodium  hydroxide  together 
with  several  small  bits  of 
phosphorus  are  placed  in 
the  flask  A,  and  a  current 
of  coal  gas  is  passed  into 
the  flask  through  the  tube 
B  until  all  the  air  has  been 
displaced.  The  gas  is  then 
turned  off  and  the  flask  FIG.  140.  The  preparation  of  phosphine 
is  heated.  Phosphine  is 

formed  in  small  quantities  and  escapes  through  the  delivery 
tube,  the  exit  of  which  is  just  covered  by  the  water  in  the 
vessel  C.  Each  bubble  of  the  gas  as  it  escapes  into  the  air 


352    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

takes  fire,  and  the  product  of  the  combustion  (P206)  forms 
beautiful  rings,  which  float  unbroken  for  a  considerable  time 
in  quiet  air.  The  pure  phosphine  does  not  take  fire  sponta- 
neously. When  prepared  as  directed  above,  a  second  hydride 
of  phosphorus  (P2H4)  is  present  which  imparts  this  property. 

Oxides  of  phosphorus.  Phosphorus  forms  two  well-known 
oxides  —  the  trioxide  (P2O3)  and  the  pentoxide  (PgO6), 
sometimes  called  phosphoric  anhydride.  When  phosphorus 
burns  in  an  insufficient  supply  of  air  the  product  is  partially 
the  trioxide ;  in  oxygen  or  an  excess  of  air  the  pentoxide 
is  formed.  The  pentoxide  is  much  the  better  known  of 
the  two.  It  is  a  snow-white  voluminous  powder  whose 
most  marked  property  is  its  great  attraction  for  water. 
It  has  no  chemical  action  upon  most  gases,  so  that  they 
can  be  very  thoroughly  dried  by  allowing  them  to  pass 
through  properly  arranged  vessels  containing  phosphorus 
pentoxide. 

Acids  of  phosphorus.  The  important  acids  of  phosphorus 
are  the  following: 

H8PO3 phosphorous  acid 

H3PO4 phosphoric  acid 

H4P2O7 pyrophosphoric  acid 

HPO3 metaphosphoric  acid 

These  may  be  regarded  as  combinations  of  the  oxides  of 
phosphorus  with  water  according  to  the  equations  given  in 
the  discussion  of  the  characteristics  of  the  family  (p.  347). 
1.  Phosphorous  acid  (H3PO^).  Neither  the  acid  nor  its 
salts  are  at  all  frequently  met  with  in  chemical  operations. 
It  can  be  easily  obtained,  however,  in  the  form  of  trans- 
parent crystals  when  phosphorus  trichloride  is  treated  with 
water  and  the  resulting  solution  evaporated: 

PC18  +  3  H2O  — >-  H3PO3  +  3  HC1 


THE  PHOSPHORUS  FAMILY  353 

It  is  a  powerful  reducing  agent  because  of  its  tendency 
to  take  up  oxygen  and  pass  over  into  phosphoric  acid. 

2.  Orthophosphoric  acid  (phosphoric  acid)  (H3POJ.  This 
acid  can  be  obtained  by  dissolving  phosphorus  pentoxide 
in  boiling  water,  as  represented  in  the  equation 


It  is  usually  made  by  treating  calcium  phosphate  with 
concentrated  sulfuric  acid.  The  calcium  sulfate  produced 
in  the  reaction  is  nearly  insoluble  and  can  be  filtered 
off,  leaving  the  phosphoric  acid  in  solution.  Very  pure 
acid  is  made  by  oxidizing  phosphorus  with  nitric  acid. 
It  forms  large  colorless  crystals  which  are  exceedingly 
soluble  in  water. 

Orthophosphates.  Since  phosphoric  acid  is  a  tribasic 
acid,  it  forms  acid  as  well  as  normal  salts.  Thus  the 
following  compounds  of  sodium  are  known: 

NaH2PO4    ......     sodium  dihydrogen  phosphate 

Na2HPO4    ......     disodium  hydrogen  phosphate 

Na3PO4  .......     normal  sodium  phosphate 

These  salts  may  be  prepared  by  bringing  together  phos- 
phoric acid  and  appropriate  quantities  of  sodium  hydrox- 
ide. Phosphoric  acid  also  forms  mixed  salts,  that  is,  salts 
containing  two  different  metals.  The  most  familiar  com- 
pound of  this  kind  is  microcosmic  salt,  which  has  the 
formula  Na(NH4)HPO4. 

The  orthophosphates  form  an  important  class  of  salts. 
The  normal  salts  are  nearly  all  insoluble  and  many  of 
them  occur  in  nature.  The  monohydrogen  phosphates  are 
as  a  rule  insoluble,  while  most  of  the  dihydrogen  salts 
are  soluble. 

The  rock  phosphates  of  Florida  and  Tennessee  contain 
about  seventy  per  cent  of  calcium  phosphate  Cag(PO4)a. 


354    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

They  are  mined  in  large  quantities  and  used  in  the  manu- 
facture of  fertilizers  (Fig.  141). 

3.  Pyrophosphoric  acid  (fl4P207).  On  heating  orthophosphoric 
acid  to  about  225°,  pyrophosphoric  acid  is  formed  in  accordance 
with  the  following  equation : 

2  H3P04 *  H4P207  +  H20 

It  is  a  white  crystalline  solid.    Its  salts  can  be  prepared  by 
heating  a  monohydrogen  phosphate : 

2  Na2HP04 *  Na4P207  +  H20 


FIG.  141.   Mining  phosphate  rock  in  Florida 

4.  Metaphosphoric  acid  (glacial  phosphoric  acid)  (HP03).    This 
acid  is  formed  when  orthophosphoric  acid  is  heated  above  400°: 

HP0 


It  is  also  formed  when  phosphorus  pentoxide  is  treated  with 
cold  water:  p^  +  H>0  —  *  2  HPO, 

It  is  a  white  crystalline  solid,  and  is  so  stable  towards  heat 
that  it  can  be  fused  and  even  volatilized  without  decomposi- 
tion. On  cooling  from  the  fused  state  it  forms  a  glassy  solid, 
and  on  this  account  is  often  called  glacial  phosphoric  acid.  It 


THE  PHOSPHOKUS  FAMILY  355 

possesses  the  property  of  dissolving  small  quantities  of  metallic 
oxides,  with  the  formation  of  compounds  which,  in  the  case  of 
certain  metals,  have  characteristic  colors.  It  is  therefore  used 
in  the  detection  of  these  metals. 

While  the  monohydrogen  phosphates,  on  heating,  give  salts 
of  pyrophosphoric  acid,  the  dihydrogen  phosphates  yield  salts 
of  metaphosphoric  acid.  The  equations  representing  these 
reactions  are  as  follows : 

2  Na2HP04 >•  Na4P207  +  H20, 

'  NaH2P04 >-  NaP08  +  H20 

Sulfides  of  phosphorus.  A  number  of  compounds  consist- 
ing of  phosphorus  and  sulfur  can  be  obtained  by  heating 
the  two  elements  in  various  proportions.  The  most  impor- 
tant of  these  has  the  formula  P4S3  and  is  called  phosphorus 
sesquisulfide. 

Chlorides  of  phosphorus.  Phosphorus  burns  readily  in 
chlorine  and  forms  the  liquid  trichloride,  PC18,  or  the 
solid  pentachloride,  PC15,  according  to  the  amount  of 
chlorine  available.  Both  of  these  compounds  have  impor- 
tant uses  in  the  preparation  of  many  organic  compounds. 

ARSENIC 

Occurrence.  Arsenic  occurs  in  considerable  quantities 
in  nature  as  the  native  element,  as  the  sulfides  realgar 
(As2S2)  and  orpiment  (As2Sg),  as  oxide  (As2O8),  and  as 
a  constituent  of  many  metallic  sulfides,  such  as  arseno- 
pyrite  (FeAsS). 

Preparation.  The  element  is  prepared  by  purifying  the 
native  arsenic  or  by  heating  the  arsenopyrite  in  iron 
tubes,  out  of  contact  with  air.  In  the  latter  case  the 
reaction  is  expressed  by  the  following  equation : 

4FeAsS 


356    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  arsenic,  being  volatile,  condenses  in  chambers  con- 
nected with  the  heated  tubes.  It  is  also  made  from  the 
oxide  by  reduction  with  carbon: 

2  As2O8  +  3  C  — *  As4  +  3  CO2 

Properties  and  conduct.  Arsenic  is  a  steel-gray,  metallic- 
looking  substance  of  density  5.73.  Like  phosphorus,  its 
molecules  contain  four  atoms  (As4).  Though  resembling 
metals  in  appearance,  it  is  quite  brittle,  being  easily 
powdered  in  a  mortar.  When  strongly  heated  it  sublimes ; 
that  is,  it  passes  into  a  vapor  without  melting,  and  con- 
denses again  to  a  crystalline  solid  when  the  vapor  is 
cooled.  Like  phosphorus  it  can  be  obtained  in  several 
allotropic  forms.  It  alloys  readily  with  some  of  the  metals 
and  finds  its  chief  use  as  an  alloy  with  lead  which  is  used 
for  making  shot,  the  alloy  being  harder  than  pure  lead 
and  forming  rounder  drops  (shot).  When  heated  on  char- 
coal with  the  blowpipe,  arsenic  is  converted  into  an  oxide 
which  volatilizes,  leaving  the  charcoal  unstained  by  any 
oxide  coating.  It  burns  readily  in  chlorine  gas,  forming 
arsenic  trichloride: 

As4  +  6Cl2 — )-4AsCl8 

Unlike  most  of  its  compounds,  the  element  itself  is  not 
poisonous. 

Arsine  (AsH3).  When  any  compound  containing  arsenic 
is  brought  into  the  presence  of  nascent  hydrogen,  arsine 
(AsH3),  corresponding  to  phosphine  and  ammonia,  is 
formed.  The  reaction  when  oxide  of  arsenic  is  so  treated  is 

As2O8  + 12  [H]  — >•  2  AsH3  +  3  H2O 

Arsine  is  a  gas  with  a  peculiar  garlic-like  odor  and  is 
intensely  poisonous.  A  single  bubble  of  the  pure  gas  has 


THE  PHOSPHORUS  FAMILY 


357 


been  known  to  prove  fatal.  It  is  an  unstable  compound, 
decomposing  into  its  elements  when  heated  to  a  moderate 
temperature.  It  is  combustible  and  burns  with  a  pale 
bluish-white  flame  to  form  arsenic  trioxide  and  water  when 
air  is  in  excess  : 

2  AsH3  +  3  O2  —  *  As2O3  +  3  H2O 

When  the  supply  of  air  is  deficient,  water  and  metallic 
arsenic  are  formed:- 

4AsH 


These  reactions  make  the  detection  of  even  minute  quan- 
tities of  arsenic  a  very  easy  problem. 

Marsh's  test  for  arsenic.  The  method  devised  by  Marsh  for 
detecting  arsenic  is  most  frequently  used.  The  apparatus  is 
shown  in  Fig.  142.  Hydrogen  is  generated  in  the  flask  A  by 
the  action  of  dilute  ,-, 

sulfuric  acid  on  zinc, 
is  dried  by  being 
passed  over  calcium 
chloride  in  the  tube 
B,  and  after  passing 
through  the  hard- 
glass  tube  C,  is  ig- 
nited at  the  jet  D. 
If  a  substance  con- 

taining  arsenic  is  FlG  142  Marsh's  apparatus  for  the  detection 
now  introduced  into  .  :  Of  arsenic 

the  generator  A,  the 

arsenic  is  converted  into  arsine  by  the  action  of  the  nascent 
hydrogen  and  passes  to  the  jet  along  with  the  hydrogen.  If 
the  tube  C  is  strongly  heated  at  some  point  near  the  middle, 
the  arsine  is  decomposed  while  passing  this  point,  and  the 
arsenic  is  deposited  just  beyond  the  heated  point  in  the  form 
of  a  shining  brownish-black  mirror.  A  small  fraction  of  a 


358    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

milligram  of  arsenic  can  be  detected  by  this  test.  If  the  tube 
is  not  heated,  the  arsine  burns  along  with  the  hydrogen  at  the 
jet.  Under  these  conditions  a  small  porcelain  dish  crowded 
down  into  the  flame  is  blackened  by  a  spot  of  metallic  arsenic, 
for  the  arsine  is  decomposed  by  the  heat  of  the  flame,  and 
the  arsenic,  cooled  below  its  kindling  temperature  by  the  cold 
porcelain,  deposits  upon  it  as  a  black  spot  (p.  317). 

Oxides  of  arsenic.  Arsenic  forms  two  oxides,  As2O3  and 
As2O5,  corresponding  to  the  oxides  of  phosphorus.  Of 
these  arsenious  oxide,  or  arsenic  trioxide  (As2O3),  is  much 
the  better  known  and  is  the  substance  usually  called  white 
arsenic,  or  merely  arsenic.  It  is  found  as  a  mineral,  but 
is  more  often  obtained  as  a  by-product  in  various  indus- 
tries in  which  metallic  sulfides  are  burned  in  air.  The 
sulfides  contain  small  amounts  of  arsenic,  and  when  they 
are  burned,  arsenious  oxide  is  formed  as  a  vapor  together 
with  sulfur  dioxide : 

2  FeAsS  +  5  02 *  Fe2O3  +  As2O3  +  2  SO2 

The  arsenious  oxide  is  condensed  in  appropriate  chambers. 
It  is  obtained  either  as  a  white  crystalline  powder  or  in 
large  vitreous  lumps  resembling  lumps  of  porcelain  in 
appearance.  It  is  extremely  poisonous,  from  0.2  to  0.3  g. 
being  a  fatal  dose.  It  is  frequently  given  as  a  poison, 
since  it  is  nearly  tasteless  and  does  not  act  rapidly.  This 
slow  action  is  due  to  the  fact  that  it  is  not  very  soluble 
and  hence  is  absorbed  slowly  by  the  system.  Arsenious 
oxide  is  also  used  as  a  chemical  reagent  in  glass  making, 
in  the  dye  industry,  and  in  the  manufacture  of  arsenical 
insecticides. 

Acids  of  arsenic.  Like  the  corresponding  oxides  of  phospho- 
rus, the  oxides  of  arsenic  are  acid  anhydrides.  In  solution  they 
combine  with  bases  to  form  salts  corresponding  to  the  salts  of 


THE  PHOSPHOEUS  FAMILY  359 

the  acids  of  phosphorus.    Thus,  we  have  salts  of  the  following 

acids : 

H8As08 arsenious  acid 

H8As04 orthoarsenic  acid 

H4As207 pyroarsenic  acid 

HAs08 metarsenic  acid 

Several  other  acids  of  arsenic  are  also  known.  Not  all  of 
these  can  be  obtained  as  free  acids,  since  they  tend  to  lose 
water  and  form  the  oxides.  Thus,  instead  of  obtaining  arse- 
nious acid  (H8As08),  the  oxide  As303  is  obtained : 

2  H8As08 >•  As308  +  3  H20 

Salts  of  all  the  acids  are  known,  however,  and  some  of  them 
have  commercial  value.  Most  of  them  are  insoluble,  and  some 
of  the  copper  salts,  which  are  green,  are  used  as  pigments. 
Paris  green,  which  has  a  complicated  formula,  is  a  well-known 
insecticide.  Lead  arsenate,  whose  formula  is  somewhat  vari- 
able, is  extensively  used  as  a  spray  on  fruit  trees. 

Sulfides  of  arsenic.  When  hydrogen  sulfide  is  passed  into 
an  acidified  solution  containing  an  arsenic  compound  the 
arsenic  is  precipitated  as  a  bright  yellow  sulfide,  thus : 

2  H8As08  +.  3  H2S +  As2S8  +  6  H20 

2  H8As04  +  5  HaS >  As2S6  +  8  H20 

In  this  respect  arsenic  resembles  the  metallic  elements, 
many  of  which  produce  sulfides  under  similar  conditions. 
The  sulfides  of  arsenic,  both  those  produced  artificially  and 
those  found  in  nature,  are  used  as  yellow  pigments. 

ANTIMONY 

Occurrence.  Antimony  occurs  in  nature  chiefly  as  the 
sulfide  (Sb2S3),  called  stibnite,  though  it  is  also  found  as 
the  oxide  and  as  a  constituent  of  many  complex  minerals. 

Preparation  and  properties.  Antimony  is  prepared  from 
the  sulfide  in  a  very  simple  manner.  The  sulfide  is 
melted  with  scrap  iron  in  a  furnace,  when  the  iron 


360    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

combines  with  the  sulfur  to  form  a  liquid  layer  of  melted 
iron  sulfide,  while  the  heavier  liquid,  antimony,  settles 
to  the  bottom  and  is  drawn  off  from  time  to  time.  The 
reaction  involved  is  represented  by  the  equation 

Sb2S8  +  3  Fe  — >•  2  Sb  +  3  FeS 

Antimony  is  a  bluish-white  metallic-looking  substance 
whose  density  is  6.52.  It  is  highly  crystalline,  hard,  and 
very  brittle.  It  has  a  rather  low  melting  point  (630°) 
and  expands  noticeably  on  solidifying. 

Chemical  conduct.  In  chemical  properties  antimony  re- 
sembles arsenic  in  many  particulars.  It  forms  the  oxides 
Sb2O3  and  Sb2O5,  and  in  addition  Sb2O4.  It  combines  with 
the  halogen  elements  with  great  energy,  burning  brilliantly 
in  chlorine  to  form  antimony  trichloride  (SbCl3).  When 
heated  on  charcoal  with  the  blowpipe  it  is  oxidized  and 
forms  on  the  charcoal  a  coating  of  antimony  oxide  which 
has  a  characteristic  bluish-white  color. 

Stibine  (SbH3).  The  gas  stibine  (SbH3)  is  formed 
under  conditions  which  are  very  similar  to  those  which 
produce  arsine,  and  it  closely  resembles  the  latter  com- 
pound, though  it  is  still  less  stable.  It  is  very  poisonous. 

In  Marsh's  test  for  arsenic  any  antimony  that  is  present 
is  converted  into  stibine,  and  this  results  in  a  black  mirror 
deposit,  as  in  the  case  of  arsenic.  The  deposit  is  more  sooty  in 
appearance  than  is  the  arsenic  deposit,  and  it  is  not  dissolved 
by  a  solution  of  sodium  hypochlorite,  whereas  the  deposited 
arsenic  is  dissolved  by  this  reagent. 

Acids  of  antimony.  The  oxides  Sb208  and  Sb208  are  weak  acid 
anhydrides  and  are  capable  of  forming  two  series  of  acids  cor- 
responding in  formulas  to  the  acids  of  phosphorus  and  arsenic. 
They  are  much  weaker,  however,  and  are  of  little  practical 
importance. 


THE  PHOSPHORUS  FAMILY  361 

Sulfides  of  antimony.  Antimony  resembles  arsenic  in  that  hy- 
drogen sulfide  precipitates  it  as  a  sulfide  when  the  hydrogen  sul- 
fide  is  conducted  into  an  acid  solution  containing  an  antimony 
compound : 

2  SbCl8  +  3  H2S »-  Sb2S3  +  6  HC1 

2  SbCl5  +  5  H2S >•  Sb2S5  +  10  HC1 

The  two  sulfides  of  antimony  are  called  the  trisulfide  and 
the  pentasulfide  respectively.  When  prepared  in  this  way 
they  are  orange-colored  substances,  though  the  mineral  stibnite 
is  black.  The  sulfides  of  antimony  are  used  in  the  manufacture 
of  matches  and  of  red  rubber. 

Metallic  properties  of  antimony.  The  physical  properties 
of  the  element  are  those  of  a  metal,  and  the  fact  that  its 
sulfide  is  precipitated  by  hydrogen  sulfide  shows  that  it 
acts  like  a  metal  in  a  chemical  way.  Many  other  reac- 
tions show  that  antimony  has  more  of  the  properties  of 
a  metal  than  of  a  nonmetal.  The  hydroxide,  Sb(OH)3, 
corresponding  to  arsenious  acid,  while  able  to  act  as  a 
weak  acid,  is  also  able  to  act  as  a  weak  base  with  strong 
acids.  For  example,  when  it  is  treated  with  concentrated 
hydrochloric  acid,  antimony  chloride  is  formed: 

Sb(OH)8  +  3HCl — >-SbCl8  +  3H2O 

The  hydroxides  of  a  number  of  elements  resemble  the 
hydroxide  of  antimony  in  that  they  can  act  either  as  a 
base  or  as  an  acid.  They  are  called  amphoteric  hydroxides. 

Hydrolysis  of  antimony  salts.  Antimony  hydroxide, 
Sb(OH)3,  is  a  very  weak  base,  and  we  should  expect  its 
salts  to  be  decomposed,  or  hydrolyzed,  by  water  (p.  226). 
If  antimony  chloride  were  to  be  completely  hydrolyzed, 
the  equation  would  be  as  follows: 

/Cl  /OH 

Sb-Cl  +  3  H20 »-Sb-OH  +  3  HC1 

XC1  VOH 


362    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  reaction  is  not  so  complete,  however,  only  two  of  the 
three  chlorine  ions  being  replaced  by  hydroxyl  ions: 

/Cl  /OH 

Sb-Cl  +  2H2O VSb-OH  +  2  HC1 

XC1  XC1 

If  we  wish  to  prevent  this  hydrolysis,  we  must  add  hydro- 
chloric acid  in  sufficient  quantity  to  reverse  the  reaction 
of  hydrolysis  by  mass  action  (p.  220). 

Basic  salts  and  oxysalts.  The  compound  formed  by 
the  partial  hydrolysis  of  antimony  chloride  is  unlike  any 
we  have  yet  met.  Since  it  contains  hydroxyl  radicals 
combined  with  a  metal,  we  must  regard  it  as  a  base ;  but 
it  also  contains  a  chlorine  atom  combined  with  a  metal, 
so  that  it  is  likewise  a  salt.  Since  it  has  the  character- 
istics of  both  a  base  and  a  salt,  it  is  a  basic  salt  (p.  187). 

/OH  /Cl 

Sb-OH  +  HC1 >- Sb-OH  +  H2O 

XOH  XOH 

base  acid  basic  salt      water 

The  basic  chloride  of  antimony  easily  loses  water,  as 
shown  in  the  equation 

/OH 
Sb-OH  : HSbQfj  (or  SbO  •  Cl) 

01 

The  resulting  compound  is  at  once  both  an  oxide  and  a 
salt.  It  is  called  antimony  oxychloride. 

BISMUTH 

Occurrence.  Bismuth  is  usually  found  in  the  uncom- 
bined  form  in  nature.  It  also  occurs  as  oxide  and  sulfide. 
Most  of  the  bismuth  of  commerce  comes  from  Saxony, 
from  Mexico,  and  from  Colorado,  but  it  is  not  an  abundant 
element. 


THE  PHOSPHORUS  FAMILY        •        363 

Preparation  and  properties.  Bismuth  is  prepared  by 
merely  heating  the  ore  containing  the  native  bismuth  and 
allowing  the  melted  metal  to  run  out  into  suitable  ves- 
sels. Other  ores  are  converted  into  oxides  and  reduced 
by  heating  with  carbon. 

Bismuth  is  a  heavy,  crystalline,  brittle  metal  nearly 
the  color  of  silver,  but  with  a  slightly  rosy  tint  which 
distinguishes  it  from  other  metals.  It  melts  at  a  low 
temperature  (271°)  and  has  a  density  of  9.8.  It  is  not 
acted  upon  by  the  air  at  ordinary  temperatures. 

Chemical  conduct.  When  heated  with  the  blowpipe  on 
charcoal,  bismuth  gives  a  coating  of  the  oxide  Bi2O3. 
This  has  a  yellowish-brown  color  which  easily  distin- 
guishes it  from  the  oxides  formed  by  other  metals.  It. 
combines  very  readily  with  the  halogen  elements,  pow 
dered  bismuth  burning  readily  in  chlorine.  It  is  below 
hydrogen  in  the  electrochemical  series  (p.  191),  and  in  the 
absence  of  air  it  is  not  acted  upon  by  hydrochloric  acid. 

Compounds  of  bismuth.  Unlike  the  other  elements  of 
this  group,  bismuth  has  almost  no  acid  properties.  Its 
chief  oxide,  Bi2O3,  is  basic  in  its  properties.  It  dissolves 
in  strong  acids  and  forms  salts  of  bismuth : 

Bi2O3  +  6  HC1  — >•  2  BiCl8  +  3  H2O 
Bi208  +  6  HN  08  — »  2  Bi(N08)8  +  3  H2O 
The  nitrate  and  the  chloride  of  bismuth  can  be  obtained 
as  well-formed  colorless  crystals. 

Bismuth  hydroxide  is  a  weak  base,  and  its  salts,  like 
those  derived  from  antimony  hydroxide,  undergo  partial 
hydrolysis  in  dilute  solution.  Thus,  bismuth  chloride  is 
hydrolyzed  according  to  the  equation 

/Cl  /OH 

Bi-Cl  +  2  II2O:z=±:Bi-OH  +  2  HC1 
^Cl  XC1 


364    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

This  action  is  reversible,  and  by  adding  an  excess  of 
hydrochloric  acid  the  basic  chloride  is  changed  into  the 
normal  chloride. 

By  loss  of  water  the  basic  chloride  'is  converted  into 
bismuth  oxy chloride : 

Bi(OH)2Cl >•  BiO  •  Cl  +  H2O 

In  a  similar  way  bismuth  nitrate  forms  the  compound 
BiO  •  NO3,  known  as  bismuth  oxynitrate,  or  subnitrate. 

ALLOYS 

Some  metals  when  melted  together  thoroughly  inter- 
mix, and  on  cooling  form  a  metallic-appearing  substance 
called  an  alloy.  Not  all  metals  will  mix  in  this  way,  and 
in  some  cases  definite  chemical  compounds  are  formed 
and  separate  out  as  the  mixture  solidifies,  thus  destroy- 
ing the  uniform  quality  of  the  alloy.  In  general,  the 
melting  point  of  the  alloy  is  below  the  average  of  the 
melting  points  of  its  constituents,  and  it  is  usually  lower 
than  that  of  any  one  of  them. 

Alloys  of  antimony  and  bismuth.  Both  antimony  and 
bismuth  readily  alloy  with  many  other  metals.  The  alloys 
so  formed  are  heavy,  are  easily  melted,  do  not  oxidize 
easily  nor  act  upon  water,  and,  in  general,  are  well 
adapted  to  many  technical  uses.  The  manufacture  of 
alloys  constitutes  the  chief  use  of  these  two  metals. 

Antimony  imparts  to  its  alloys  the  property  of  expand- 
ing slightly  in  solidification,  which  renders  them  especially 
useful  in  type  founding,  where  fine  lines  are  to  be  repro- 
duced on  a  cast.  Type  metal  consists  of  antimony,  lead, 
and  tin.  Babbitt  metal,  used  for  journal  bearings  in 
machinery,  contains  the  same  metals  in  a  different  pro- 
portion, together  with  a  small  percentage  of  copper. 


THE  PHOSPHOKUS  FAMILY  365 

Bismuth  is  particularly  valuable  in  the  production  of 
very  low-melting  alloys.  For  example,  Wood's  metal,  con- 
sisting of  bismuth,  lead,  tin,  and  cadmium,  melts  at  60.5°. 
The  low  melting  point  of  such  alloys  is  turned  to  practical 


FIG.  143.   An  automatic  fire  curtain  above  a  door 

account  in  making  automatic  fire  curtains  and  automatic 
water  sprinklers  in  buildings,  safety  plugs  in  boilers,  and 
many  similar  devices. 

Fig.  143  shows  a  fire  curtain,  which  is  held  in  place  by  two 
wires  (A,  A)  joined  at  B  by  a  bismuth  alloy.  In  case  of  fire  the 
alloy  melts,  and  the  wires  holding  the  curtain  up  are  thereby 
released  and  the  curtain  drops,  covering  the  door. 

EXERCISES 

1.  What  is  the  derivation  of  the  word  phosphorus  ? 

2.  What  compounds  would  you  expect  phosphorus  to  form  with 
bromine  and  iodine?    Write  the  equations  showing  the  action  of 
water  on  these  compounds. 

3.  In  the  preparation  of  phosphine  why  is  coal  gas  passed  into 
the  flask?    What  other  gases  would  serve  the  same  purpose? 

4.  Give  the   formula  for  the  salt  which  phosphine  forms  with 
hydriodic  acid.    Give  the  name  of  the  compound. 

5.  Could  phosphoric  acid  he  substituted  for  sulfuric  acid  in  the 
preparation  of  the  common  acids? 


366    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

6.  Write  the  equations  for  the  preparation  of  the  three  sodium 
salts  of  orthophosphoric  acid. 

7.  Why  does  a  solution  of  normal  sodium  phosphate  show  an 
alkaline  reaction  ? 

8.  On  the  supposition  that  bone  ash  is  pure  calcium  phosphate, 
what  weight  of  it  would  be  required  in  the  preparation  of  1  kg.  of 
phosphorus  ? 

9.  If  arsenopyrite  is  heated  in  a  current  of  air,  what  products 
are  formed? 

10.  (a)  Write  equations  for  the  complete  combustion  of  hydro- 
gen sulfide,  methane,  and   arsine.     (6)  In  what  respects  are  the 
reactions  similar? 

11.  Write  the  equations  for  all  the  reactions  involved  in  Marsh's 
test  for  arsenic. 

12.  Write  the  names  and  formulas  for  the  acids  of  antimony. 

13.  W'rite  the  equations  showing  the  hydrolysis  of  antimony  tri- 
chloride ;  of  bismuth  nitrate. 

14.  In  what  respects  does  nitrogen  resemble  the  members  of  the 
phosphorus  family  ? 

15.  What  weight  of  arsenic  trioxide  can  be  prepared  from  1  kg. 
of  arsenopyrite? 

16.  Suppose  you  wish  to  prepare  1kg.  of  bismuth  oxychloride; 
what  weight  of  bismuth  would  be  required  ? 

17.  What  weight  of  stibnite  is  necessary  for  the  preparation  of 
10  kg.  of  antimony  ? 


CHAPTER  XXIX 
SILICON;  TITANIUM;  BORON 


NAME  OF  ELEMENT 

SYMBOL 

ATOMIC 

WEIGHT 

DENSITY 

CHLORIDES 

OXIDES 

Silicon      .     . 

Si 

28  3 

2  35 

SiCl 

SiO 

Titanium  
Boron  

Ti 
B 

48.1 
11.0 

4.5 
2.45 

TiCl4 
BCU 

TiO2 
B  (X 

General.  Each  of  the  'three  elements  silicon,  titanium, 
and  boron  belongs  to  a  separate  periodic  family,  but  they 
occur  near  together  in  the  periodic  grouping  and  are  very 
similar  both  in  properties  and  in  chemical  activity.  Since 
the  other  elements  in  their  families  either  are  so  rare  that 
they  need  not  be  studied  in  detail  or  are  best  understood 
in  connection  with  other  elements,  it  is  convenient  to 
consider  these  three  together  at  this  point. 


SILICON 

Occurrence.  Next  to  oxygen,  silicon  is  the  most  abun- 
dant element,  for  the  solid  crust  of  the  earth  is  estimated 
to  contain  "28  per  cent  of  this  element.  All  varieties  of 
granite,  gneiss,  sandstone,  shale,  clay,  and  marl  contain 
large  percentages  of  silicon  —  limestone  and  dolomite  being 
the  only  important  geological  formations  measurably  free 
from  it.  To  some  extent  its  compounds  are  assimilated  by 
plants  and  animals,  and  silicon  compounds  constitute  the 
outer  shell  of  many  aquatic  organisms. 
367 


368    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  element.   In  the  laboratory  crystallized  silicon  is  best 
prepared  by  the  reduction  of  the  dioxide  with  aluminium  : 

3Si0 


The  silicon  dissolves  in  the  excess  of  melted  aluminium, 
and  when  the  solution  has  cooled  and  become  solid,  the 
aluminium  is  dissolved  in  hydrochloric  acid,  the  silicon 
being  left  in  the  form  of  shining  metallic  needles.  The 
reduction  of  the  dioxide  with  carbon  has  always  presented 
the  difficulty  that  the  reduced  element  tends  to  combine 
with  excess  of  carbon  to  form  a  carbide.  This  difficulty 
has  been  overcome  to  a  great  extent,  and  nearly  pure  sili- 
con is  now  manufactured  in  large  quantities.  By  reducing 
a  mixture  of  the  oxides  of  silicon  and  iron  with  carbon, 
an  alloy  of  the  two  elements,  called  ferrosilicon,  is  obtained. 
This  alloy,  as  well  as  the  purer  silicon,  finds  an  important 
application  in  the  metallurgy  of  iron. 

Properties.  The  element  presents  a  close  analogy  with 
carbon  in  that  it  can  be  obtained  in  amorphous  form,  as 
well  as  in  crystals  resembling  the  diamond.  These  crystals 
are  very  hard,  easily  scratching  glass,  and  have  a  density 
of  2.35.  Jhey  melt  at  about  1420°.  A  lump  of  the  ele- 
ment is  very  brittle  and  breaks  with  a  crystalline  fracture 
which  has  a  metallic,  silvery  appearance. 

At  ordinary  temperatures  silicon  is  inactive.  At  high  tem- 
peratures it  combines  with  most  elements,  forming  silicides, 
such  as  those  of  magnesium  (Mg2Si)  and  carbon  (CSi). 

Compounds  of  silicon  with  hydrogen  and  the  halogens. 
Silicon  hydride  (SiH4)  corresponds  in  formula  to  methane 
(CH4),  but  its  properties  are  more  like  those  of  phosphine 
(PHS).  It  is  a  very  inflammable  gas  of  disagreeable  odor 
and,  as  ordinarily  prepared,  takes  fire  spontaneously  on 
account  of  the  presence  of  impurities. 


SILICON;  TITANIUM;  BORON  369 

Silicon  combines  with  the  elements  of  the  chlorine 
family  to  form  such  compounds  as  SiCl4  and  SiF4.  'Of 
these  silicon  fluoride  (SiF4)  is  the  most  familiar  and  inter- 
esting. As  stated  in  the  discussion  of  fluorine,  it  is  formed 
when  hydrofluoric  acid  acts  on  silicon  dioxide  or  on  a 
silicate.  With  silica  the  reaction  is  thus  expressed  : 

SiO2  +  2  HaF2  —  >•  SiF4  +  2  H2O 

Silicon  fluoride  is  a  very  volatile,  invisible,  poisonous  gas. 
In  contact  with  water  it  is  partially  decomposed,  as  shown 
in  the  equation 

SiF4  +  4  H2O  —  v-  2  H2F2  +  Si(OH)4 

The  hydrofluoric  acid  so  formed  combines  with  an  addi- 
tional amount  of  silicon  fluoride,  forming  the  complex 
fluosilicic  acid  (H2SiF6),  thus: 

HF  +  SiF 


Silicides.  As  the  name  indicates,  silicides  are  compounds 
consisting  of  silicon  and  some  one  other  element.  They  are 
very  stable  at  high  temperatures  and  are  usually  made  by 
heating  the  appropriate  substances  in  an  electric  furnace. 

The  most  important  silicide  is  carborundum,  which  is  a 
silicide  of  carbon  of  the  formula  CSi.  It  is  made  by  heat- 
ing coke  and  sand  in  an  electric  furnace,  the  process  being 
extensively  carried  on  at  Niagara  Falls.  The  following 
equation  represents  the  reaction: 


The  substance  so  prepared  consists  of  beautiful  purplish- 
black  crystals,  which  are  surpassed  in  hardness  only  by 
a  few  substances,  such  as  the  diamond  and  boron  carbide. 
Carborundum  is  used  as  an  abrasive  ;  .that  is,  as  a  mate- 
rial for  grinding  and  polishing  very  hard  substances. 


370    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Fig.  144  shows  two  samples  of  the  crystalline  material, 
as  well  as  a  whetstone  and  a  grinding  wheel  prepared 
from  carborundum,  illustrating  some  of  its  many  uses. 


FIG.  144.   Crystals  of  carborundum  and  abrasive  utensils  made 
of  carborundum 

Manufacture  of  carborundum.  The  mixture  of  materials  is 
heated  in  a  large  resistance  furnace,  similar  to  the  one  em- 
ployed in  the  manufacture  of  graphite 
(p.  119).  Fig.  145  represents  a  cross 
section  of  the  furnace  after  charging,  A 
being  the  carbon  core  and  B  the  coke  and 
sand.  Fig.  146  shows  the  appearance 
after  heating.  A  is  the  core  of  carbon,  sur- 
rounded by  crystallized  carborundum  B. 
Around  this  is  a  shell  of  amorphous  carbo- 
rundum C,  while  D  is  unchanged  charge. 

Silicon  dioxide  (silica)  (Si02).  This 
substance  is  found  in  a  great  variety 
of  forms  in  nature,  both  in  the  amorphous  and  in  the 
crystalline  condition.  In  the  form  of  quartz  (Fig.  147) 


FIG.  145.    Cross  section 
of    a    charged    carbo- 
rundum furnace  before 
being  heated 


SILICON;  TITANIUM;  BORON 


371 


FIG.  146.   Cross  section 
of    a    charged    carbo- 
rundum furnace  after 
being  heated 


it  is  found  in  beautifully  formed  six-sided  prisms,  some- 
times of  great  size.  When  pure  it  is  perfectly  transparent 
and  colorless.  Some  colored  varieties 
are  given  special  names,  as  amethyst 
(violet),  rose  quartz  (pale  pink),  smoky 
quartz  (black),  milky  quartz  (white). 
Other  varieties  of  silicon  dioxide,  some 
of  which  also  contain  water,  are  chal- 
cedony, onyx,  jasper,  opal,  agate,  and 
flint.  Sand  and  sandstone  are  largely 
silicon  dioxide. 

The  skeletons  of  certain  microor- 
ganisms (infusoria)  are  composed  of  nearly  pure  silica.  In 
some  localities  these  have  accumulated  in  immense  deposits, 
forming  a  very  fine  and  sharp  sand  called  infusorial  earth. 
This  material  is  often 
used  as  a  scouring 
powder,  especially 
in  scouring  soaps. 

Properties.  As 
obtained  by  chemi- 
cal processes  silicon 
Dioxide  is  an  amor- 
phous white  powder. 
In  the  crystallized 
state  it  is  very  hard 
and  has  a  density 
of  2.6.  Pure  silica 
begins  to  soften  at 
about  1600°,  and 
somewhat  above  this  temperature  it  can  be  drawn  out  into 
threads,  blown  like  glass  into  tubes  and  small  vessels,  and 
molded  into  large  bowls  and  pipes  for  use  in  chemical 


FIG.  147.   A  cluster  of  quartz  crystals 


372    AN  ELEMENTARY  STUDY  OF  CHEMISTKY 


industries.  These  articles  are  attacked  by  comparatively  few 
ordinary  reagents,  and  they  do  not  expand  or  contract  to  any 
appreciable  extent  with  even  very  great  changes  in  temper- 
ature. On  this  account  a  quartz  vessel  can  be  heated  red-hot 
and  plunged  into  cold  water  without  cracking.  Fig.  148 

shows  a  quartz  crucible  and 
quartz  tubes  on  a  wire  tri- 
angle used  to  support  the 
crucible  when  it  is  heated. 
Chemical  conduct.  Silica 
is  insoluble  in  water  and  in 
most  acids.  It  is  very  stable, 
so  that  the  oxygen  which  it 
contains  can  be  removed  only 
by  the  most  powerful  reducing 
agents  and  at  very  high  tem- 
peratures. Hydrofluoric  acid 

FIG.  148.  A  crucible  and  a  triangle      attacks  k  readi1^    (P«  267>' 
made  from  quartz  according  to  the  equation 


Si02  +  2H2F2 


SiF4  +  2H20 


Since  it  is  the  anhydride  of  an  acid,  it  dissolves  in  fused 
alkalies  to  form  silicates.  Being  nonvolatile,  it  will  drive 
out  most  other  anhydrides  when  it  is  heated  with  their 
salts  to  a  high  temperature,  especially  when  the  silicates 
so  formed  are  fusible.  The  following  equations  illustrate 
this  property: 

Na2CO8  +  SiO2  — >-  Na2Si03  +  CO2 
Na2SO4  +  SiO2  — >-'Na2SiO8  +  SO8 

Simple  silicic  acids.  Silicon  forms  two  simple  acids, 
orthosilicic  acid  (H4SiO4)  and  metasilicic  acid  (H2SiO3). 
Orthosilicic  acid  is  set  free  as  a  jellylike  mass  when 


SILICON;  TITANIUM;  BORON  373 

orthosilicates  are  treated  with  strong  acids.   If  one  attempts 
to  dry  this  acid,  it  loses  water,  passing  into  metasilicic  acid  : 

HSi0 


Metasilicic  acid,  when  heated,  breaks  up  into  silica  and 
water,  thus: 


Both  of  these  silicic  acids  are  very  weak,  and  their  soluble 
salts  are  much  hydrolyzed  in  solution. 

Salts  of  silicic  acids  ;  silicates.  A  number  of  salts  of  the 
orthosilicic  and  metasilicic  acids  occur  in  nature.  Thus, 
mica  (KAlSiO4)  is  a  mixed  salt  of  orthosilicic  acid,  and 
wollastonite  (CaSiOg)  is  a  metasilicate. 

Condensed  silicic  acids.  Silicon  has  the  power  to  form 
a  great  many  complex  acids  which  may  be  regarded  as 
derived  from  the  union  of  several  molecules  of  orthosilicic 
acid,  with  the  loss  of  water.  These  are  called  condensed 
silicic  acids.  For  example,  we  have 

3  H4Si04  —  +  H4Si8O8  +  4  H2O 

Salts  of  these  condensed  acids  make  up  the  great  bulk 
of  the  earth's  crust.  Feldspar,  for  example,  has  the  for- 
mula KAlSi3Og  and  is  a  mixed  salt  of  the  acid  H4Si3Og, 
whose  formation  is  represented  in  the  equation  above. 
Kaolin,  or  pure  clay,  has  the  formula  H4Al2Si2O9  or,  as  it 
is  commonly  written,  Al2Si2O7  -  2  H2O.  Granite  is  com- 
posed of  crystals  of  feldspar  and  mica  cemented  together 
with  amorphous  silica. 

Water  glass.  A  concentrated  solution  of  sodium  silicate 
(Na2SiO3)  or  of  potassium  silicate  (K2SiO3)  or  of  both  is 
called  water  glass.  It  is  a  thick,  sticky  liquid  made  by 
fusing  sand  with  the  carbonate  of  sodium  or  of  potassium. 
It  is  strongly  alkaline  in  reaction,  owing  to  the  ready 


374    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

hydrolysis  of  these  salts.  It  is  used  for  the  purpose  of 
giving  a  glazed  waterproof  surface  to  porous  materials, 
such  as  wood,  stone,  and  plaster;  to  render  curtains  non- 
inflammable  ;  as  a  glue  for  glass  and  pottery  ;  and  as  an 
ingredient  in  cheap  soaps. 

Its  property  of  closing  pores  is  turned  to  account  in  the  pre- 
serving of  eggs  for  winter  use.  The  eggs  are  packed  in  crocks 
and  then  covered  with  a  liquid  made  by  adding  1  volume  of 
commercial  water  glass  to  10  volumes  of  water.  Over  the  liquid 
is  then  poured  a  little  melted  paraffin,  which  soon  hardens  and 
excludes  the  air.  Fresh  eggs  can  be  preserved  for  from  eight 
to  ten  months  in  this  way. 

Glass.  When  sodium  silicate,  calcium  silicate,  and  sili- 
con dioxide  are  heated  together  to  a  high  temperature, 
the  mixture  slowly  fuses  to  a  transparent  liquid,  which,  on 
cooling,  passes  into  the  rigid  material  called  glass.  Instead 
of  starting  with  the  silicates  of  sodium  and  ralcium  it  is 
more  convenient  and  economical  to  heat  sodium  carbonate 
(or  sulfate)  and  lime  with  an  excess  of  clean  sand,  the 
silicates  being  formed  during  the  heating  as  follows  : 

Na2C08  +  SiO2  —  >-  Na2SiO8  +  CO2 


Molding  and  blowing  glass.  The  way  in  which  the  melted 
mixture  is  handled  in  the  glass  factory  depends  upon  the  char- 
acter of  the  object  to  be  made.  Many  articles,  such  as  bottles, 
are  made  by  blowing  the  plastic  glass  into  hollow  molds  of  the 
desired  shape.  The  mold  is  opened,  a  lump  of  plastic  glass  on 
a  hollow  rod  is  lowered  into  it,  and  the  mold  is  then  closed. 
By  blowing  into  the  tube  the  glass  is  forced  into  the  shape  of 
the  mold.  The  mold  is  then  opened  (Fig.  149)  and  the  object 
lifted  out.  The  top  of  the  object  must  be  cut  off  at  the  proper 
place  and  the  sharp  edges  rounded  off  in  a  flame.  Bottles  are 
now  more  often  made  by  machinery,  in  which  process  the  neck 
is  finished  first  and  the  bottle  then  blown  by  compressed  air. 


SILICON;  TITANIUM;  BORON 


375 


FIG.  149.   Making  a  glass  vessel  in  a  mold 


Other  objects,  such  as  lamp  chimneys,  glasses,  and  beakers, 
are  revolved  while  being  blown  in  the  mold,  and  have  no  ridge 
showing  where  the  mold  closes.  Window  glass  is  made  by 

gathering  a  lump  of 
molten  glass  on  the 
end  of  a  hollow  rod 
(Fig.  150)  and  blow- 
ing this  into  the  form 
of  large  hollow  cylin- 
ders (Fig.  151)  about 
6  ft.  long  and  1^  ft. 
in  diameter.  These 
are  cut  longitudinally 
(Fig.  152)  and  are 
then  placed  in  an 
oven  and  heated  un- 
til they  soften,  when 
they  are  flattened  out  into  plates  and  cut  into  the  desired 
sizes.  Similar  cylinders  are  now  made  by  dipping  a  hollow 
tube  into  the  melted  glass  and  slowly  withdrawing  it  while 
compressed  air  is 
blown  through  the 
tube.  In  this  way  a 
very  long  cylinder  is 
formed.  Plate  glass 
is  cast  into  flat  slabs 
(Fig.  153),  which  are 
then  ground  and  pol- 
ished (Fig.  154)  to  per- 
fectly plane  surfaces. 
Varieties  of  glass. 
The  glass  made 
from  sodium  carbon- 
ate, lime,  and  sand 
is  soft  and  easily 
fusible.  If  potassium  carbonate  is  substituted  for  the  sodium 
carbonate,  the  glass  is  much  harder  and  less  easily  fused; 
increasing  the  amount  of  sand  has  somewhat  the  same  effect. 


FIG.  150.   First  step  in  making  window  glass  : 
blowing  the  ball 


376    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


FIG.  151.    Second  step  in  making  window 
.     glass  :  blowing  the  cylinders 


Potassium  glass,  of  which  Jena  glass  is  a  variety,  is  largely 
used  in  making   chemical    glassware,   because  it   resists   the 

action  of  reagents 
better  than  the  softer 
sodium  glass.  If  lead 
oxide  is  substituted 
for  the  whole  or  a 
part  of  the  lime,  the 
glass  is  very  soft  but 
has  a  high^index  of  re- 
fraction and  is  valua- 
ble for  making  optical 
instruments  and  arti- 
ficial jewels  (paste). 

Coloring  glass.  Va- 
rious substances  fused 
along  with  the  glass 
impart  characteristic 

colors.  The  amber  color  of  common  bottles  is  due  to  iron 
compounds  in  the  glass ;  in  other  cases  iron  colors  the  glass 
green.  Cobalt  com- 
pounds color  it  deep 
blue,  compounds  of 
manganese  give  it  an 
amethyst  tint,  and 
uranium  compounds 
impart  a  peculiar 
yellowish-green  color. 
Iron  is  nearly  always 
present  in  sand,  and 
this  makes  a  green 
glass  unless  an  oxi- 
dizing agent  is  used. 
The  green  color  is 
largely  removed  by 
the  addition  of  manganese  dioxide,  which  oxidizes  the  iron 
compounds  to  a  form  having  a  yellowish  tinge,  and  this  color  is 
then  neutralized  by  the  manganese,  since  the  yellow  produced 


FIG.  152.   Third  step  in  making  window 
glass :  cutting  the  cylinders 


SILICON;  TITANIUM;  BORON 


377 


by  the  iron  and  the  amethyst   produced  by  the   manganese 

are  complementary  colors,  producing  white. 

Nature  of  glass.    Glass  is  not  a  definite  chemical  compound, 

and  its  composition  varies  between  wide  limits.   Fused  glass  is 

really  a  solution  of 
various  silicates,  such 
as  those  of  calcium 
and  lead,  in  fused 
sodium  silicate  or 
potassium  silicate.  A 
certain  amount  of 
silicon  dioxide  is  also 
present.  This  solu- 
tion is  cooled  under 
such  conditions  that 
the  dissolved  sub- 
stances do  not  sepa- 


FlG. 


Casting  and  rolling  plate  glass 


rate  from  the  solvent.  The  compounds  which  are  used  to  color 
the  glass  are  sometimes  converted  into  silicates  which  then 
dissolve  in  the  glass,  giving  it  a  uniform  color.  In  other  cases, 
as  in  the  milky  glasses,  which  resemble  porcelain  in  appear- 
ance, the  color  or 
opaqueness  is  due  to 
the  finely  divided 
material  evenly  dis- 
tributed throughout 
the  glass,  but  not  dis- 
solved in  it.  Milky 
glass  is  made  by  mix- 
ing calcium  fluoride, 
tin  oxide,  or  some 
other  insoluble  sub-  FIG.  154.  Polishing  plate  glass 

stance  in  the  melted 

glass.  Selenium,  gold,  or  cuprous  oxide,  in  very  finely  divided 
form  scattered  through  glass,  gives  it  shades  of  red  (ruby  glass). 
Enamels.  The  surface  of  metal  vessels,  such  as  cooking 
utensils  and  bathtubs,  is  often  covered  by  a  kind  of  opaque 
glass  called  enamel  (granite  or  agate  ware).  This  contains 


378    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

boric  oxide  (B203)  in  place  of  some  silica,  and  oxides  of  a 
number  of  different  metals,  such  as  barium,  zinc,  or  lead,  in 
place  of  some  of  the  calcium. 

TITANIUM 

Occurrence.  Titanium  occurs  rather  sparingly  in  nature 
and  is  usually  found  in  the  form  of  the  dioxide  TiO2, 
called  rutile,  or  as  an  iron  titanite  (FeTiO3)  known  as 
ilmenite,  or  as  a  variable  constituent  of  certain  magnetic 
iron  ores. 

The  element.  The  element  can  be  obtained  by  the  re- 
duction of  the  dioxide  with  carbon  in  an  electric  furnace, 
but  prepared  in  this  way  it  always  contains  carbon  and 
usually  nitrogen.  Very  pure  specimens  have  been  prepared 
by  the  action  of  titanium  chloride  on  sodium  in  a  closed 
steel  bomb: 

TiCl4  +  4  Na *  Ti  +  4  NaCl 

When  the  element  contains  carbon  it  is  hard  and  very 
brittle.  Its  melting  point  is  above  that  of  platinum,  being 
about  1800°.  Its  specific  gravity  is  4.5. 

An  alloy  known  as  ferrotitanium  is  made  by  reducing 
iron  ores  rich  in  titanium  with  carbon.  This  alloy  is  manu- 
factured on  the  large  scale  at  Niagara  Falls  and  is  used 
in  preparing  titanium  steel.  It  contains  15.5  per  cent  of 
titanium,  1.4  per  cent  of  silicon,  7.5  per  cent  of  carbon,  and 
74.3  per  cent  of  iron. 

At  high  temperatures  titanium  shows  a  very  marked 
tendency  to  unite  with  nitrogen,  the  nitride  TiN  being  the 
product  of  this  direct  union.  The  nitride  is  therefore 
always  produced  in  any  attempt  to  prepare  titanium  in  an 
apparatus  to  which  air  has  access,  and  this  compound  was 
formerly  considered  to  be  the  element  itself.  When  iron 


SILICON;  TITANIUM;  BORON  379 

ores  containing  titanium  are  smelted,  a  substance  in  appear- 
ance resembling,  crystallized  copper  is  often  found  in  the 
slag  or  adhering  to  the  lining  of  the  furnace.  This  was 
also  at  one  time  supposed  to  be  the  metal,  but  is  now 
known  to  have  the  formula  Ti10C2Ng. 

The  compounds.  The  compounds  of  titanium  very  closely 
resemble  those  of  silicon.  The  dioxide  of  titanium,  like 
that  of  silicon,  is  an  acid  anhydride  and  forms  a  large  num- 
ber of  acids  closely  resembling  the  various  types  of  silicic 
acids.  These  are  even  weaker  than  those  of  silicon,  and  their 
salts  hydrolyze  more  readily.  Fluotitanic  acid  (H2TiF6) 
and  its  salts  are  well  known. 

BORON 

Occurrence.  Boron  occurs  in  nature  as  boric  acid  (HgBO3) 
and  in  salts  of  condensed  boric  acids,  which  usually  have 
very  complicated  formulas. 

Preparation  and  properties.  The  element  boron  is  ex- 
tremely difficult  to  prepare  in  pure  condition,  and  it  is^ 
only  known  in  an  amorphous  state.  Its  electrical  resistance 
varies  to  an  extraordinary  extent  with  changes  in  tempera- 
ture, and  this  property  promises  to  make  it  very  useful. 

Boric  acid  (H3B03).  This  compound  is  found  dissolved 
in  the  water  of  hot  springs  in  some  localities,  particularly 
in  Italy.  Being  volatile  with  steam,  boric  acid  is  present  in 
the  vapor  from  these  springs.  The  acid  is  easily  obtained 
from  these  sources  by  condensation  and  evaporation,  the 
necessary  heat  being  supplied  by  other  hot  springs. 

It  is  often  prepared  by  treating  a  strong  hot  solution  of 
borax  (Na2B4O7)  with  sulfuric  acid.  Boric  acid,  being  but 
sparingly  soluble  in  water,  crystallizes  out  on  cooling: 

Na2B407  +  5  H20  +  H2SO4  — *  Na2SO4  +  4  H8BO3 


380    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Boric  acid  crystallizes  in  pearly  flakes  which  are  slip- 
pery to  the  touch.  It  is  a  mild  antiseptic  and  is  often 
used  in  medicine,  particularly  for  eyewashes.  Its  acid 
properties  are  extremely  weak. 

Metabolic  acid  and  tetraboric  acid.  When  boric  acid  is 
gently  heated,  it  is  converted  into  metaboric  acid  (HBO2)  : 

HBO 


On  heating  metaboric  acid  to  a 
somewhat  higher  temperature,  tetra- 
boric acid  (H2B4O7)  is  formed: 

4HB02  —  ^H2B407  +  H20 

Borax.  The  sodium  salt  of  tetra- 
boric acid  has  the  formula  Na2B4O7. 
Borax  is  a  hydrate  (p.  251)  of 
this  salt  and  has  the  formula 


Na2B4O7  •  5  H2O  or 


.  10  H2O 


FIG.  155.    A  specimen  of 
*  the  mineral  colemanite 


according  to  the  temperature  of 
crystallization.  It  is  found  in  some 
arid  countries,  as  southern  Califor- 
nia and  Tibet,  but  is  now  made  commercially  from  the 
mineral  colemanite.  This  is  the  calcium  salt  of  a  complex 
boric  acid,  Ca2B6On  •  5  H2O,  and  occurs  in  large  deposits 
in  California.  When  it  is  treated  with  a  solution  of  sodium 
carbonate,  calcium  carbonate  is  precipitated,  and  borax 
crystallizes  from  the  solution.  Fig.  155  shows  a  specimen 
of  colemanite. 

When  heated,  borax  at  first  swells  up  greatly  and  then 
melts  to  a  clear  glass.  This  glass  has  the  property  of 
easily  dissolving  many  metallic  oxides,  and  on  this  account 
borax  is  used  in  brazing  for  the  purpose  of  removing  from 
the  metallic  surfaces  to  be  brazed  the  film  of  oxide  with 


SILICON;  TITANIUM;  BORON  381 

which  they  are  likely  to  be  covered.  These  oxides  often 
give  a  characteristic  color  to  the  clear  borax  glass,  and  on 
this  account  borax  beads  are  often  used  in  testing  for  the 
presence  of  metals. 

The  reason  that  metallic  oxides  dissolve  in  borax  is  that 
borax  contains  an  excess  of  acid  anhydride,  as  can  be  more 
easily  seen  if  its  formula  is  written  2  NaB02  +  B208.  The 
metallic  oxide  combines  with  this  excess  of  acid  anhydride, 
forming  a  mixed  salt  of  metaboric  acid. 

Borax  is  extensively  used  as  a  constituent  of  enamels 
and  glazes  for  both  metal  ware  and  pottery.  It  is  used 
to  soften  hard  water  for  domestic  purposes,  as  a  mild 
alkali  (like  soap),  as  an  antiseptic,  and  in  brazing. 


1.  Account  for  the  fact  that  both  silicon  and  carborundum  can 
be  made  by  heating  sand  with  carbon. 

2.  Account  for  the  fact  that  a  solution  of  borax  in  water  is 
alkaline. 

3.  What  weight  of  water  of  hydration  does  1  kg.  of  borax  contain  V 

4.  What  weight  of  borax  can  be  made  from  a  ton  of  colemanite  ? 
6.  Why   does    not    sodium    silicate    form    a   glass   suitable    for 

common   use  ? 

6.  Calculate  the  percentage  composition  of  feldspar. 

7.  Which  contains  the  greater  percentage  of  silicon  —  feldspar 
or  kaolin? 


CHAPTER  XXX 
THE  COLLOIDAL  STATE 

If  to  a  moderately  dilute  solution  of  water  glass  (sodium 
silicate)  a  little  sulfuric  or  hydrochloric  acid  is  added,  no 
very  striking  change  is  noticeable.  However,  if  the  test 
tube  is  set  aside  for  some  hours  and  is  then  examined,  it 
will  be  found  that  the  contents  of  the  tube  have  set  to 
a  clear,  stiff  jelly.  This  jelly  consists  essentially  of  water 
and  silicic  acid. 

In  a  similar  way  a  little  dried  gelatin  swells  up  and 
dissolves  in  warm  water,  but  sets  to  a  jelly  when  the 
solution  is  cooled.  A  great  many  substances,  especially 
organic  materials,  assume  the  jellylike  form  and  have 
long  been  called  colloids  —  a  word  that  means  "  glue." 

Two  states  of  colloids.  We  have  just  seen  that  the  two 
colloids  mentioned  are  capable  of  existing  in  a  condition 
that  at  first  sight  appears  to  be  ordinary  solution,  as 
well  as  in  the  state  of  jelly.  This  is  not,  however,  a 
true  solution,  but  a  state  of  such  finely  divided  material 
suspended  in  the  solvent  that  the  product  closely  re- 
sembles a  solution.  That  there  is  a  real  difference  can 
be  shown  by  the  following  method: 

If  a  beam  of  sunlight  passes  through  perfectly  clear 
air,  it  leaves  no  path  and  is  invisible  ;  but  if  it  passes 
through  dusty  air,  it  illuminates  each  dust  particle  and  thus 
makes  a  visible  path,  such  as  is  often  seen  in  a  dark  room 
when  a  sunbeam  enters  by  some  crack  or  hole  in  a  cur- 
tain. Now  in  a  perfectly  analogous  way  a  beam  of  light 


THE  COLLOIDAL  STATE 


shining  through'  a  true  solution,  such  as  sodium  chloride 
dissolved  in  water,  leaves  no  path,  because  the  particles 
of  salt  (molecules)  are  far  too  small  to  reflect  and  scatter 
light.  But  when  a  beam  of  light  passes  through  the 
clear  silicic  acid  before  it  sets  to  jelly  or  through  the 
gelatine  warmed  with  water,  it  leaves  a  bright  path, 
showing  the  presence  of  very  small  particles  in  suspension, 
not  in  real  solution  (Fig.  156).  This  clear  solution-like 
condition  is  called  a  sol  (or 
sometimes  hydrotof),  while  the 
jellylike  state  is  called  a  gel 
(or  hydrogel).  Many  substances 
can  exist  as  either  sol  or  gel, 
and  in  either  state  are  called 
colloids. 

Coagulation  of  colloids.  The 
changing  of  a  sol  into  a  gel  is 
called  coagulation.  The  white 
of  an  egg,  beaten  up  with  water, 
constitutes  a  typical  sol.  When 
this  is  heated  the  sol  turns  into 
an  opaque  white  gel  that  will 
not  again  pass  into  the  sol  form.  In  other  words  the 
coagulation  is  irreversible.  In  the  case  of  gelatine,  cooling 
produces  the  gel  and  heating  produces  the  sol,  and  the 
coagulation  can  be  reversed  as  often  as  may  be  desired. 

Sometimes  the  addition  of  an  electrolyte  (acid,  base, 
salt)  occasions  coagulation,  as  in  the  example  of  the  sol 
of  silicic  acid.  Plastic  clay  well  stirred  with  water  is 
partly  material  that  can  be  filtered  off  (suspended  matter) 
and  partly  material  that  is  too  finely  divided  to  be  filtered 
(sol).  The  addition  of  a  few  drops  of  sulfuric  acid 
quickly  precipitates  the  sol  as  gel,  and  so  too  does  the 


FIG.  156.-  A  beam  of  light  shin- 
ing through  a  colloidal  solution 


384    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

addition  of  certain  salts.  Alkalies  tend  to  make  the  sol 
condition  more  stable  and  may  even  cause  the  coagulated 
gel  to  pass  again  into  the  sol  form.  Sometimes  the  addi- 
tion of  one  sol  to  another  causes  the  coagulation  of  both 
as  gel,  while  in  other  cases  the  sol  state  is  rendered  more 
stable.  For  example,  when  colloidal  ferric  hydroxide  is 
added  to  colloidal  antimony  sulfide  or  colloidal  arsenious 
oxide,  both  are  precipitated.  On  the  other  hand,  gela- 
tin renders  many  sols  more 
stable. 

Preparation  of  colloidal  sols. 
In  a  general  way  it  may  be 
said  that  colloidal  sols  may 
be  produced  in  either  of  two 
ways: 

==  ' 1.  By  powdering  solids.    We 

FIG.  157.     The    preparation    of  ^^  ^b.  a  solid  and  by 

metallic    colloids    by    sparking  J         .  .J 

under  water  mechanical    means   powder  it 

until   it   is  so   finely   divided 

that  it  will  make  a  sol  when  mixed  with  water.  In  the 
case  of  metals  and  some  other  substances  this  may  be 
done  by  dipping  electrodes  made  of  the  substance  under 
water  and  causing  a  strong  electric  current  to  strike  an 
arc  between  the  electrodes  (Fig.  157).  Particles  of  the 
electrode  are  torn  off  by  the  current  and  remain  sus- 
pended in  the  water  as  sol.  The  sols  of  gold  are  purple, 
blue,  or  green  in  color,  depending  upon  the  size  of  the 
particles.  Those  of  platinum  are  brown. 

2.  By  imperfect  precipitation.  When  two  substances  are 
brought  together  that  are  capable  of  forming  an  insoluble 
precipitate,  in  general  we  expect  the  precipitate  to  form. 
Under  certain  circumstances,  especially  if  the  solutions 
are  very  dilute,  precipitation  does  not  take  place  but  the 


THE  COLLOIDAL  STATE  385 

expected  insoluble  precipitate  remains  suspended  as  sol. 
Thus,  when  chromium  chloride  or  chromium  sulfate  is 
treated  with  sodium  hydroxide  we  expect  the  precipi- 
tation of  insoluble  chromium  hydroxide : 

CrCls  +  3  NaOH >-  Cr(OH)8  +  3  NaCl 

If  the  solutions  are  cold  and  are  brought  together  slowly, 
no  precipitate  forms ;  but  after  a  time  the  contents  of  the 
test  tube  sets  to  a  jelly.  In  a  similar  way  nearly  any 
insoluble  material  may  be  obtained  either  as  a  sol  or  a  gel, 
though  very  strongly  crystalline  substances,  such  as  barium 
sulfate  (BaSO4),  are  difficult  to  obtain  in  these  forms. 

Colloids  and  hydrolysis.  We  have  seen  that  salts  of  a 
weak  acid  or  a  weak  base  undergo  some  hydrolysis  in 
solution  (p.  226).  In  many  cases  either  the  acid  or  the 
base  is  known  to  be  insoluble,  yet  precipitation- does  not 
always  occur ;  so  we  are  forced  to  assume  the  presence 
in  solution  of  a  compound  we  know  to  be  insoluble. 

Thus,  sodium  silicate  is  markedly  hydrolized,  as  is  shown 
by  its  strongly  alkaline  reaction,  and  silicic  acid  must  be 
produced  in  the  reaction.  We  know  it  'to  be  insoluble, 
yet  no  precipitate  forms  when  sodium  silicate  is  dissolved 
in  water.  As  another  example  we  may  take  the  violet- 
colored  salt  ferric  nitrate,  Fe(NO,)8  •  6  H2O.  If  a  crystal 
of  this  salt  is  placed  in  water,  it  forms  a  yellow  or  brown 
solution  that  is  strongly  acid  in  reaction,  indicating  con- 
siderable hydrolysis.  Ferric  hydroxide,  which  should  be 
formed  in  the  reaction,  is  highly  insoluble  and  should 
therefore  form  as  a  precipitate,  but  does  not. 

These  apparently  contradictory  facts  may  be  reconciled 
if  we  suppose  that  the  silicic  acid  and  the  ferric  hydroxide 
remain  as  colloidal  sols  instead  of  separating  as  gels  or 
precipitates.  Examination  by  transmitted  light  shows  that 


386    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

this  explanation  is  warranted  by  the  appearance  of  the 
hydrolized  solutions. 

Nature  of  colloids.  From  what  has  been  said  it  will  be 
seen  that  colloids  are  not  a  special  kind  of  substance,  but 
•that  the  colloidal  condition  is  one  which  any  substance 
may  assume  just  as  it  may  occur  as  a  solid,  as  a  liquid,  as 
a  gas,  or  in  true  solution.  In  this  state  it  is  too  coarsely 
divided  to  be  a  true  solution,  yet  too  finely  divided  to  be 
removed  by  filtration  or  settling.  It  will  also  be  apparent 
that  there  can  be  no  very  sharp  distinction  drawn  between 
a  colloid  and  a  fine  suspension,  the  one  passing  by  imper- 
ceptible stages  into  the  other. 

It  is  probable  that  each  particle  of  a  typical  colloid 
consists  of  thousands  of  molecules,  but  that  these  are 
clumped  together  without  any  special  order.  When  the 
particles  assume  an  orderly  arrangement  they  constitute 
crystals,  and  crystalline  particles  continue  to  grow  in  size, 
precipitating  from  the  solution  in  definite  solid  form. 
On  the  other  hand,  the  colloid  clumps  tend  to  form  net- 
works or  spongelike  forms  that  inclose  water  and  consti- 
tute jellies. 

Colloids  and  the  industries.  Many  important  industries 
depend  upon  the  properties  of  colloids.  Many  colored 
glasses,  such  as  ruby  glass,  owe  their  color  to  colloidal 
material  dispersed  through  the  glass.  Photographic  plates 
are  made  of  colloidal  salts  of  silver,  dispersed  through 
gelatin  or  some  similar  substance.  Pastes  of  various 
kinds,  bread  dough,  glues,  and  cements  owe  many  of 
their  properties  to  their  colloidal  character.  Soils  con- 
tain a  great  deal  of  colloidal  material,  both  organic  and 
mineral,  and  their  ability  to  hold  water  and  to  support 
the  growth  of  plants  is  greatly  dependent  upon  this 
fact.  Rubber  is  a  typical  colloid,  as  are  other  similar 


THE  COLLOIDAL  STATE 


387 


plastic  materials,  such  as  gums  and  waxes.  Many  min- 
erals have  originated  from  the  gradual  hardening  of  gels 
and  owe  their  peculiar  structure  to  this  origin,  as,  for 
example,  the  banded  structure  of  onyx  and  agate.  Indeed, 
we  are  just  beginning  to  realize  how  much  of  the  chemistry 
of  everyday  life  is  in  the  realm 
of  colloidal  materials. 

Emulsions.  If  water  and  a 
little  kerosene  are  vigorously 
shaken  together,  the  resulting 
liquid  is  milky  in  appearance 
and  consists  of  very  small  drops 
of  kerosene  scattered  through 
the  water.  These  drops  rapidly 
run  together  and  collect  into 
a  layer  of  oil  floating  on  the 
water,  and  the  milky  appear- 
ance entirely  vanishes.  If  the 
experiment  is  repeated,  but 
with  the  addition  of  a  very 
small  particle  of  soap  before 
the  shaking,  the  product  re- 
mains milky  for  a  long  time 
(Fig.  158).  Indeed,  it  may  re- 
main so  for  months.  Such  a 
milky  liquid  is  called  an  emul- 
sion. Under  other  conditions 

water  and  kerosene  produce  an  emulsion  which  consists  of 
droplets  of  water  dispersed  in  kerosene.  Milk  is  a  natural 
emulsion,  but  not  a  very  perfect  one,  for  the  cream  rises 
to  the  top  rather  rapidly. 

It  will  be  seen  that  the  essential  distinction  between  a 
colloid  and  an  emulsion  is  that  in  the  first  case  the  particles 


FIG.  158.   Emulsions 

In  A  the  oil  floats  on  the  water; 

in  B  an  emulsion  has  been  formed 

by  adding  soap 


388    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

are  those  of  a  solid,  while  in  the  case  of  an  emulsion 
they  are  liquid  drops.  We  may  expect  any  one  liquid  to 
form  an  emulsion  with  any  other,  provided  the  two  do 
not  mix  to  form  a  solution  as  do  water  and  alcohol.  For 
an  emulsion  to  be  at  all  stable  it  appears  to  be  necessary 
that  a  third  colloidal  substance  should  be  present,  such  as 
the  soap  in  the  experiment  with  kerosene,  or  the  casein 
of  milk. 

Fogs.  If  the  air  is  perfectly  free  from  solid  particles,  it 
may  become  greatly  supersaturated  with  water  vapor.  If 
now  dust  or  smoke  mingles  with  the  supersaturated  air, 
each  little  dust  particle  becomes  a  center  upon  which 
water  vapor  condenses,  and  since  the  solid  particles  are 
very  numerous,  the  water  droplets  are  also  very  numer- 
ous and  very  small.  The  result  of  such  condensation  is 
what  we  call  a  fog.  It  will  be  seen  that  it  is  but  natural 
that  large  cities,  especially  those  situated  near  large  bodies 
of  water  or  rivers  and  in  humid  regions,  should  be  subject 
to  fog.  From  a  chemical  standpoint  a  fog  might  be  con- 
sidered as  an  emulsion  in  which  the  water  droplets  are 
dispersed  through  a  gas  instead  of  through  another  liquid. 


CHAPTER  XXXI 
THE  METALS 

The  metals.  The  elements  so  far  considered  have  nearly 
all  been  those  whose  compounds  with  oxygen  and  hydrogen 
are  adds,  and  they  are  called  the  acid-forming  elements 
or  the  nonmetals.  Those  which  we  shall  now  study  are 
known  collectively  as  the  metals.  The  hydroxides  of  the 
metals  are  bases,  and  on  this  account  the  metals  are  some- 
times defined  as  those  elements  whose  hydroxides  are  bases. 
When  the  hydroxide  of  a  metal  or  any  of  the  simple  salts 
derived  from  the  hydroxide  are  dissolved  in  water,  the 
metallic  element  forms  the  cation  and  carries  a  positive 
charge. 

The  distinction  between  a  metal  and  a  nonmetal  is  not 
a  very  sharp  one,  since  the  hydroxides  of  a  number  of 
elements  act  as  bases  under  some  conditions  and  as  acids 
under  others.  We  have  seen  that  antimony  is  an  element 
of  this  kind. 

Properties  of  the  metals.  The  metals  are  all  solids  except 
mercury,  which  is  a  liquid.  Most  metals  have  a  high  den- 
sity, although  three  of  them,  namely,  lithium,  sodium,  and 
potassium,  are  lighter  than  water.  As  a  rule,  the  metals 
are  good  conductors  of  heat  and  electricity,  are  notably 
crystalline  in  structure,  and  take  a  bright  polish.  With 
the  exception  of  gold  and  copper,  they  have  a  silvery 
luster.  Most  of  them  combine  readily  with  oxygen  and 
sulfur,  and  their  surfaces  quickly  tarnish  on  exposure  to 


390    AN  ELEMENTARY  STUDY  OF  CHEMISTEY 

the  air.  A  few  of  them,  such  as  gold  and  platinum,  have 
little  chemical  activity  and  so  retain  their  luster  in  the  air. 

Occurrence  of  metals  in  nature.  A  few  of  the  metals  are 
found  in  nature  in  the  free  state.  Among  these  are  gold, 
platinum,  and  frequently  copper.  They  are  usually  found 
combined  with  other  elements  in  the  form  of  oxides  or 
salts  of  various  acids.  Silicates,  carbonates,  sulfides,  and 
sulfates  are  the  most  abundant  of  these  salts,  All  inor- 
ganic substances  occurring  in  nature,  whether  they  con- 
tain a  metal  or  not,  are  called  minerals.  Those  minerals 
from  which  a  useful  substance  can  be  extracted  are  called 
ores  of  the  substance. 

Extraction  of  metals:  metallurgy.  The  process  of  ex- 
tracting a  metal  from  its  ores  is  called  the  metallurgy  of 
the  metal.  The  metallurgy  of  each  metal  presents  pecu- 
liarities of  its  own,  but  there  are  several  methods  of  general 
application  which  are  very  frequently  employed. 

1.  Reduction  of  an  oxide  with  carbon.  Many  of  the  metals 
occur  in  nature  in  the  form  of  oxides.  When  some  of 
these  oxides  are  heated  to  a  high  temperature  with  carbon, 
the  oxygen  combines  with  the  carbon  and  the  metal  is  set 
free.  Iron,  for  example,  occurs  largely  in  the  form  of  the 
oxide  Fe2O3.  When  this  is  heated  with  carbon  the  reaction 
expressed  in  the  following  equation  takes  place: 


FeaO8-{-3C 

Many  ores  other  than  oxides  may  be  changed  into  oxides 
which  can  then  be  reduced  by  carbon.  The  conversion  of 
such  ores  into  oxides  is  generally  accomplished  by  heating, 
and  the  process  is  called  roasting.  Many  carbonates  and 
hydroxides  decompose  directly  into  the  oxide  on  being 
heated.  Sulfides,  on  the  other  hand,  must  be  heated  in  a 
current  of  air,  the  oxygen  of  the  air  entering  into  the 


THE  METALS  391 

reaction.    The  following  equations  will  serve  to  illustrate 
these  changes  in  the  case  of  the  ores  of  iron: 


FeCO3 

2  Fe(OH)3  —  >•  Fe2O3  +  3  H2O 
4  FeS2  +  11  O2  —  *  2  Fe203  +  8  SO2 

2.  Reduction  of  an  oxide  with  aluminium.    Not  all  oxides, 
however,  can  be  reduced  by  carbon.     In  such  cases  alu- 
minium may  be  used.    Thus,  chromium  may  be  obtained 
in  accordance  with  the  following  equation  : 

Cr203  +  2  Al  -  *  2  Cr  +  A12O8 

This  method  was  first  used  by  the  German  chemist  Gold- 
schmidt  and  is  called  the  Goldschmidt  method. 

3.  Electrolysis.    In  recent  years  increasing  use  is  being 
made  of  the  electric  current  in  the  preparation  of  metals. 
In  some  cases  the  separation  of  the  metal  from  its  com- 
pounds is  accomplished  by  passing  the  current  through  a 
solution  of  a  suitable  salt  of  the  metal,  the  metal  usually 
being  deposited   upon   the  cathode.     In   other  cases  the 
current  is  passed  through  a  fused  salt  of  the  metal,  the 
chloride  being  best  adapted  to  this  purpose. 

Electrochemical  industries.  Most  of  the  electrochemical 
industries  of  the  country  are  carried  on  where  water  power 
is  abundant,  since  this  furnishes  the  cheapest  means  for 
the  generation  of  electrical  energy.  Niagara  Falls  is  the 
most  important  locality  in  this  country  for  such  industries, 
and  many  different  electrochemical  products  are  manu- 
factured there  (Fig.  159).  Some  industries  depend  upon 
electrolytic  processes,  while  in  others  the  electrical  energy 
is  used  merely  as  a  source  of  heat  In  the  latter  case  the 
process  is  called  thermoelectric. 


392    AN  ELEMENTARY  STUDY  OF  CHEMISTKY 

Preparation  of  compounds  of  the  metals.  Since  the  com- 
pounds of  the  metals  are  so -numerous  and  so  varied  in 
character^  there  are  many  ways  of  preparing  them.  In 
many  cases  the  properties  of  the  substance  te  be  prepared, 
or  the  material  available  for  its  preparation,  suggest  a  rather 
unusual  way.  There  are,  however,  a  number  of  general 
principles  which  are  constantly  applied  in  the  preparation 
of  the  compounds  of  the  metals,  and  a  clear  understanding 


FIG.  159.    Electrochemical  power  plants  at  Niagara  Falls 

of  them  will  save  much  time  and  effort  in  remembering 
the  details  in  any  given  case.  Some  of  the  general  methods 
for  the  preparation  of  compounds  are  the  following: 

1.  By  direct  union  of  two  elements.  This  is  usually  ac- 
complished by  heating  the  two  elements  together.  Thus, 
the  sulfides,  chlorides,  and  oxides  of  a  metal  can  generally 
be  obtained  in  this  way.  The  following  equations,  serve 
as  examples  of  this  method: 

Fe  +  S — >-FeS 
Cu  +  Cl2 
2Mg  +  O2 


THE  METALS  393 

2.  By  the  decomposition  of  a  compound.  This  decomposition 
may  be  brought  about  either  by  heat  alone  or  by  the  com- 
bined action  of  heat  and  a  reducing  agent.  Thus,  when  the 
nitrate  of  a  metal  is  heated  the  oxide  of  the  metal  is  usu- 
ally obtained.  Copper  nitrate,  for  example,  decomposes  as 
follows : 


-  ^  2  CuO  +  4  NO2  +  O2 
Similarly,  the  carbonates  of  the  metals  yield  oxides,  thus  : 


CaCO3 

Most  of  the  hydroxides  form  an  oxide  and  water  when 
heated  :  g  A1(OH)>  -  ^  A1>O|  +  3  Uf) 

When  heated  with  carbon,  sulfates  are  reduced  to  sulfides, 
BaSO4  +  2  C  —  *  BaS  +  2  CO2 

3.  Methods  based  on  equilibrium  in  solution.    In  the  prep- 
aration of  compounds  the  first  requisite  is  that  the  reactions 
chosen  shall  be  of  such  a  kind  as  will  go  on  to  completion. 
In  the  chapter  on  chemical  equilibrium  (p.  224)  it  was 
shown  that  reactions  in  solution  may  become  complete  in 
either  of  three  ways:    (1)  a  gas  may  be  formed  which 
escapes  from  solution  ;  (2)  an  insoluble  solid  may  be  formed 
which  precipitates  ;  (3)  two  different  ions  may  combine  to 
form  undissociated  molecules.    By  the  judicious  selection 
of  materials  these  principles  may  be  applied  to  the  prepa- 
ration of  a  great  variety  of.  compounds,  and  illustrations 
of   such  methods  will  very  frequently  be  found  in   the 
subsequent  pages. 

4.  By  fusion  methods.    It  sometimes  happens  that  sub- 
stances which  are  insoluble  in  water  and  in  acids,  and 
which  cannot  therefore  be  brought  into  double  decomposi- 
tion in  the  usual  way,  are  soluble  in  other  liquids,  and 


394    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

when  dissolved  in  them  can  be  decomposed  and  converted 
into  other  desired  compounds.  Thus,  barium  sulfate  is 
not  soluble  in  water,  and  sulfuric  acid,  being  less  volatile 
than  most  other  acids,  cannot  easily  be  driven  out  from 
this  salt.  When  brought  into  contact  with  melted  sodium 
carbonate,  however,  it  dissolves  in  it  and,  since  barium 
carbonate  is  insoluble  in  melted  sodium  carbonate,  double 
decomposition  takes  place : 

Na2COg  +  BaSO4 >-  BaCO3  -f-  Na2SO4 

When  the  cooled  mixture  is  dissolved  in  water  the  sodium 
sulfate  formed  in  the  reaction,  together  with  any  excess  of 
sodium  carbonate  which  may  be  present,  dissolves.  The 
barium  carbonate  can  then  be  filtered  off  and  converted 
into  any  desired  salt  by  the,  processes  already  described. 

Important  insoluble  compounds.  Since  precipitates  play 
so  important  a  part  in  the  reactions  which  substances 
undergo,  as  well  as  in  the  preparation  of  many  chemical 
compounds,  it  is  important  to  know  what  substances  are 
insoluble  in  water.  Knowing  this,  we  can  in  many  cases 
predict  reactions  under  certain  conditions,  and  are  assisted 
in  devising  ways  to  prepare  desired  compounds.  While 
there  is  no  general  rule  which  will  enable  one  to  foretell 
the  solubility  of  any  given  compound,  nevertheless  a  few 
general  statements  can  be  made  which  will  be  of  much 
assistance. 

1.  Hydroxides.    All  hydroxides    are  insoluble  in   water 
save    those    of    ammonium,    sodium,   potassium,   calcium, 
barium,  and  strontium. 

2.  Nitrates.    All  nitrates  are  soluble  in  water. 

3.  Chlorides.    All  chlorides  are   soluble    in  water  save 
silver  and   mercurous   chlorides.     (Lead   chloride  is  but 
slightly  soluble.) 


THE  METALS  395 

4.  Sulfates.    All  sulfates  are  soluble  in  water  save  those 
of  barium,  strontium,  and  lead.    (Sulfates  of  silver  and  cal- 
cium are  only  moderately  soluble.) 

5.  Sulfides.    All   sulfides    are   insoluble   in   water  save 
those  of  ammonium,  sodium,  and  potassium.    The  sulfides 
of  calcium,  barium,  strontium,  and  magnesium  are  insoluble 
in  water,  but  are  changed  by  hydrolysis  into  acid  sulfides 
which  are  soluble ;   on  this  account  they  cannot  be  pre- 
pared by  precipitation. 

6.  Carbonates,  phosphates,  and  silicates.    All  normal  car- 
bonates, phosphates,  and  silicates  are  insoluble  in  water 
save  those  of  ammonium,  sodium,  and  potassium. 

EXERCISES 

1 .  Write  equations  representing  four  different  ways  for  preparing 
Cu(N03)2. 

2.  Write  equations  representing  six  different  ways  for  preparing 
ZnSO4. 

3.  Write  equations  for  two  reactions  to  illustrate  each  of  the 
three  ways  in  which  reactions  in  solutions  may  become  complete. 

4.  Give  one  or  more  methods  for  preparing  each  of  the  following 
compounds:    CaCl2,  PbCl2,  BaSO4,  CaCO8,  (NH4)2S,  Ag2S,  PbO, 
Cu(OH)2  (for  solubilities,  see  last  paragraph  of  chapter).    State  in 
each  case  the  general  principle  involved  in  the  method  of  preparation 
chosen. 


CHAPTER  XXXII 
THE  ALKALI  METALS 


METAL 

ATOMIC 

MELTIXG 

Lithium  (Li)  .     .     . 

6.94 

0.534 

186.0° 

Arfvedson,  1817 

Sodium  (Na)  .     .     . 

23:00 

0.971 

97.5° 

Davy,  1807 

Potassium  (K)    .     . 

39.10 

0.862 

62.3° 

Davy,  1807 

Rubidium  (Rb)  .     . 

85.45 

1.532 

38.0° 

Bunsen,  1861 

Caesium  (Cs)  .     .     . 

132.81 

1.87 

26.4° 

Bunsen,  1860 

Characteristics  of  the  family.  The  elements  listed  in 
the  above  table  constitute  a  family  in  Group  I  of  the 
periodic  table.  They  are  called  the  alkali  metals  for  the 
reason  that  the  most  familiar  members  of  the  family, 
namely,  sodium  and  potassium,  are  constituents  of  com- 
pounds that  have  long  been  known  as  alkalies.  Before 
considering  each  element  separately  it  is  advisable  to 
discuss  briefly  the  family  as  a  whole  since  a  number  of 
statements  can  be  made  that  apply  to  all  of  these  elements. 

1.  Occurrence.     While  none  of  these  metals  occur  free 
in  nature,  their  compounds  are  widely  distributed,  being 
found  in   sea  and  mineral  waters,  in  salt  beds,  and  in 
many  rocks.    Sodium  and  potassium  are,  however,  the  only 
ones  that  occur  in  abundance. 

2.  Preparation.    The  alkali  metals  are  most  readily  pre- 
pared by   the   electrolysis    of   their  fused   hydroxides   or 
chlorides.    They  may  also  be  prepared  by  the  reduction 
of  their  oxides,  hydroxides,  or  carbonates. 


THE  ALKALI  METALS 


397 


3.  Properties.    They  are  soft  metals,  easily  molded  by 
the  fingers.     They  have  low   melting   points   and   small 
densities,  as  shown  in  the  table.    Their  densities  (sodium 
excepted)  are  in  the  same  order  as  their  atomic  weights, 
while  their  melting  points 

and  boiling  points  are  in 
the  reverse  order.  The  pure 
metals  have  a  silvery  luster, 
but  tarnish  at  once  when  ex- 
posed to  the  air,  because  of 
the  formation  of  a  film  of 
oxide  upon  their  surface : 
hence  they  are  often  pre- 
served in  some  liquid,  such 
as  kerosene,  which  contains 
no  oxygen.  They  stand 
at  the  head  of  the  electro- 
chemical series  of  the  metals 
(p.  191)  and  in  general  are 
very  active  elements. 

4.  Compounds.    The  alkali      FIG.  160.    Robert  Wilhelm  Bunsen 
metals  act  as  univalent  ele-  (1811-1899) 

ments       in      the      formation       A  distinguished  German  chemist  who 

.  discovered  rubidium  and  caesium,  m- 

OI     Compounds.      Their    hy-       vented  the  spectroscope  and  the  lab- 

droxides  (MOH)  are  white    'oratolT  burner'  and  contributed  to 

v  many  advances  in  .chemistry 

solids  and  are  very  soluble 

in  water.  In  dilute  aqueous  solutions  these  hydroxides  are 
largely  ionized  and  to  about  the  same  extent,  forming  the 
ions  M+  and  OH~  ;  hence  their  solutions  are  strongly  basic. 
With  few  exceptions  the  salts  of  the  alkali  metals  are 
white  solids,  and,  unless  otherwise  stated,  it  will  be  so 
understood  in  the  description  of  the  individual  compounds. 
With  the  exception  of  lithium  these  metals  form  very  few 


398    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

insoluble  compounds,  so  that  it  is  difficult  to  prepare  their 
compounds  by  precipitation.  The  compounds  of  sodium 
and  potassium  are  so  similar  in  properties  that  for  most 
purposes  they  can  be  used  interchangeably.  Those  of 
sodium  are  cheaper  than  the  corresponding  compounds  of 
potassium  and  so  are  more  largely  used. 

Sodium  and  potassium  are  the  only  important  members 
of  the  family.  Sodium,  as  well  as  its  hydroxide,  has  been 
described  in  Chapter  XV,  and  this  description  should  be 
carefully  reviewed  in  connection  with  the  present  chapter. 

COMPOUNDS  OF  SODIUM 

General.  In  addition  to  sodium  hydroxide  (p.  175) 
sodium  forms  a  large  number  of  compounds.  With  the 
exception  of  the  nitrate  all  these  compounds  are  prepared 
from  the  chloride,  since  it  is  so  abundant  and  inexpensive. 
The  processes  involved  are  often  complicated,  owing  to 
the  fact  that  the  compounds  of  sodium  are  all  soluble 
and  therefore  cannot  be  prepared  directly  from  the  chlo- 
ride by  precipitation ;  moreover,  the  chloride  is  a  salt  of 
a  strong  acid  and  is  not  readily  acted  upon  by  most  other 
acids.  Experiments  have  shown  that  the  most  economical 
method  of  procedure  consists  either  in  first  changing  the 
chloride  into  the  hydroxide'  by  the  electrolysis  of  its 
aqueous  solution  or  in  converting  it  into  the  carbonate 
by  the  methods  to  be  described.  Since  the  hydroxide  is  a 
base  and  the  carbonate  is  a  salt  of  a  very  volatile  acid, 
both  are  readily  changed  into  other  compounds. 

Sodium  peroxide  (Na20a).  Since  sodium  is  a  univalent 
element,  we  should  expect  it  to  form  an  oxide  of  the  formula 
Na2O.  While  such  an  oxide  can  be  prepared,  the  peroxide, 
Na2O2,  is  much  better  known.  It  is  a  yellowish-white 


THE  ALKALI  METALS  399 

powder  made  by  burning  sodium  in  air.  Its  chief  use  is 
as  an  oxidizing  agent.  When  heated  with  oxidizable  sub- 
stances it  gives  up  a  part  of  its  oxygen,  as  shown  in  the 
equation 


2Na202 

Water  decomposes  it  in  accordance  with  the  equation 
Na202  +  2  H2O  —  *  2  NaOH  +  H2O2 

Acids  act  readily  upon  it,  forming  a  sodium  salt  and 
hydrogen  peroxide  : 

Na202  +  2  HC1  —  »  2  NaCl  +  H2O2 

In  these  last  two  reactions  the  hydrogen  dioxide  formed 
decomposes  into  water  and  oxygen  unless  precautions  are 
taken  to  prevent  the  temperature  from  rising  (p.  27)  : 

2H202  —  ^2H20  +  02 

Sodium  chloride  (common  salt)  (NaCl).  Sodium  chloride, 
or  common  salt,  is  very  widely  distributed  in  nature. 
Thick  strata,  evidently  deposited  by  the  evaporation  of 
salt  water,  are  found  in  many  places.  In  the  United 
States  the  most  important  localities  for  salt  are  New 
York,  Michigan,  Ohio,  and  Kansas.  Sometimes  the  salt  is 
mined,  especially  if  it  is  in  the  pure  form  called  rock  salt. 
More  frequently  a  strong  brine  is  pumped  from  deep  wells 
sunk  into  the  salt  deposit.  The  brine  is  evaporated  either 
by  beating  or,  in  the  preparation  of  the  coarser  grades  of 
salt,  by  simply  exposing  the  brine  to  the  air  (Fig.  161). 
Salt  crystallizes  in  the  form  of  small  cubes. 

Uses  of  salt.  Since  salt  is  so  abundant  in  nature,  it  is 
used  in  the  preparation  of  nearly  all  compounds  containing 
either  sodium  or  chlorine.  These  include  many  substances 
of  the  highest  importance  to  civilization,  such  as  soap,  glass, 


400    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

hydrochloric  acid,  soda,  and  bleaching  powder.  Enormous 
quantities  of  salt  are  therefore  produced  each  year.  Small 
quantities  are  essential  to  the  life  of  man  and  animals. 
Pure  salt  does  not  absorb  moisture ;  the  fact  that  ordinary 
salt  becomes  moist  when  exposed  to  air  is  not  due  to  a 
property  of  the  salt  but  to  impurities  occurring  in  it,  espe- 
cially to  the  presence  of  calcium  and  magnesium  chlorides. 


FIG.  161.   The  evaporation  of  salt  brine  in  the  open  air 

Sodium  bromide  (NaBr) ;  sodium  iodide  (Nal).  These  com- 
pounds resemble  sodium  chloride  in  their  physical  properties. 
They  can  be  prepared  by  the  action  of  bromine  and  iodine  re- 
spectively upon  a  solution  of  sodium  hydroxide  (p.  409).  They 
are  used  to  a  limited  extent  in  photography  and  in  medicines. 

Sodium  sulfate  (Na2SOJ.  This  salt  is  prepared  by  the 
action  of  sulfuric  acid  upon  sodium  chloride,  hydrogen 
chloride  being  formed  at  the  same  time  (p.  167): 

2  NaCl  +  H2SO4  — ->  Na2SO4  +  2  HC1 

It  is  also  prepared  by  the  action  of  sodium  chloride  upon 
magnesium  sulfate,  the  latter  being  obtained  in  large 
quantities  in  the  manufacture  of  potassium  chloride: 

MgSO4  +  2  NaCl  — >-  Na2SO4  +  MgCl2 


THE  ALKALI  METALS  401 

The  anhydrous  sodium  sulfate  is  a  white  solid.  It  is 
readily  soluble  in  water  and,  under  ordinary  conditions, 
crystallizes  out  as  the  hydrate,  Na2SO4  •  10  H2O  (known 
as  G-lauber's  salt).  Large  quantities  of  sodium  sulfate  are 
used  in  making  sodium  carbonate  and  glass.  The  salt  is 
also  used  in  medicine. 

Sulfites  of  sodium.  The  acid  sulfite,  NaHS08,  often  called 
sodium  bisulfite,  is  formed  by  saturating  a  solution  of  sodium 
carbonate  with  sulfur  dioxide.  Sulfurous  acid  is  first  formed 
by  the  union  of  the  dioxide  with  water,  and  this  decomposes 
the  carbonate : 

Na2C03  +  2  H2S08 >•  2  NaHS03  +  C03  +  H20 

The  normal  sulfite,  NaaSOg,  is  prepared  by  adding  sodium 
carbonate  to  a  saturated  solution  of  the  acid  sulfite  in  the 
proportion  indicated  in  the  following  equation : 

2  NaHSO,  +  NaaC08 >-  2  Na2S03  +  H20  +  C02 

Both  of  the  sulfites  of  sodium  readily  absorb  oxygen,  form- 
ing the  corresponding  sulfates  ;  they  are  therefore  reducing 
agents.  They  are  used  to  some  extent  as  bleaching  agents 
and  as  preservatives. 

Sodium  thiosulfate  (Na2S203).  This  salt  is  made  by 
boiling  a  solution  of  sodium  sulfite  with  sulfur: 

Na2SO8+S — ^Na,S8O, 

The  hydrate,  Na2S2O3  •  5  H2O,  is  frequently  called  sodium 
hyposulfite,  or  simply  hypo.  It  is  used  in  photography 
and  in  the  bleaching  industry  to  absorb  the  excess  of 
chlorine  which  is  left  upon  the  bleached  fabrics. 

Sodium  carbonate  (Na2C03).  There  are  two  different 
methods  now  employed  in  the  manufacture  of  this  salt. 

1.  Leblanc  process.  This  older  process  involves  several 
distinct  reactions,  as  shown  in  the  following  equations: 


402    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 
(a)  Sodium  chloride  is  first  converted  into  sodium  sulfate: 
2  NaCl  +  H2S04  —  *•  Na2SO4  +  2  HC1 

(6)  The  sodium  sulfate  is  next  reduced  to  sodium  sul- 
fide  by  heating  it  with  carbon: 


Na2S04  +  2  C  —  ^Na2S  +  2  CO2 

(<?)  The  sodium  sulfide  is  then  heated  with  calcium  car- 
bonate, when  the  following  reaction  takes  place  : 

Na2S  +  CaCO3  -  >-CaS  +  Na2CO3 

2.  Solvay  process.  This  more  modern  process  depends 
upon  reactions  taking  place  in  solution  and  represented 
in  the  equations 

NaCl  +  NH4HC03  —  >-  NaHCO3  +  NH4C1        (1) 
2  NaHCO3  —  *•  Na2C03  +  H2O  +  CO2  (2) 

When  concentrated  solutions  of  sodium  chloride  and 
of  ammonium  hydrogen  carbonate  are  brought  together, 
the  sparingly  soluble  sodium  hydrogen  carbonate  is  pre- 
cipitated as  represented  in  equation  (1).  This,  by  heat- 
ing, is  converted  into  the  normal  carbonate  as  indicated 
in  equation  (2).  The  ammonium  chloride  formed  (equa- 
tion (1))  is  treated  with  lime  (p.  202),  ammonia  being 
liberated.  This  ammonia  together  with  water  and  the 
carbon  dioxide  generated  as  indicated  in  equation  (2) 
combine  to  form  ammonium  hydrogen  carbonate: 

NH3  +  C02  +  H20  —  >•  NH4HC03 

This  is  treated  with  salt,  and  the  process  is  begun  over 
again,  according  to  equation  (1). 

Historical.  In  former  times  sodium  carbonate  was  made  by 
burning  seaweeds  and  extracting  the  carbonate  from  their  ash. 


THE  ALKALI  METALS 


403 


On  this  account  the  salt  was  called  soda  ash,  and  the  name  is 
still  in  common  use.  During  the  French  Eevolution  this  supply 
was  cut  off,  and  in  behalf  of  the  French  government  Leblanc 
(Fig.  162)  made  a  study  of 
methods  of  preparing  the  car- 
bonate directly  from  common 
salt.  As  a  result  he  devised  the 
method  which  bears  his  name 
and  which  was  used  exclusively 
for  many  years.  Although  the 
method  is  still  used  in  England, 
it  has  been  entirely  replaced  in 
the  United  States  by  the  Solvay 
process,  which  was  devised  by  the 
Belgian  chemist  Solvay  in  1863. 
By-products.  The  substances 
obtained  in  a  given  process, 
aside  from  the  main  product,  are 
called  the  by-products.  The  suc- 
cess of  many  processes  depends 
upon  the  value  of  the  by-products 
formed.  Thus,  hydrochloric  acid, 
a  by-product  in  the  Leblanc  proc- 
ess, is  valuable  enough  to  make 
the  process  pay,  even  though 
sodium  carbonate  can  be  made 
more  cheaply  in  other  ways. 

Properties  of  sodium  carbon- 
ate. Sodium  carbonate  forms 
large  crystals  of  the  formula 
Na2CO8  •  10  H2O,  known  as 
washing  soda,  or  sal  soda.  The 
monohydrate,  Na2CO3  .  H2O, 
is  also  prepared  and  used  commercially  to  some  extent. 

An  aqueous  solution  of  sodium  carbonate  has  a  mild 
alkaline  reaction  and  is  used  for  laundry  purposes.    Mere 


FIG.  162.   Nicolas  Leblanc 
(1742-1806) 

Inventor  of  the  first  method  of  pre- 
paring sodium  carbonate  from  salt. 
(Statue  erected  in  Paris) 


404    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

mention  of  the  fact  that  the  salt  is  used  in  the  manu- 
facture of  glass,  soap,  and  many  chemical  reagents  will 
indicate  its  importance  in  the  industries.  It  is  one  of 
the  few  soluble  carbonates. 

Sodium  hydrogen  carbonate  (NaHC03).    This  salt,  called 
bicarbonate  of 'soda,  or  baking  soda,  is  made  by  the  Solvay 


Fiu.  163.   A  deposit  of  sodium  nitrate  in  Chile 

process,  as  explained  above,  or  by  passing  carbon  dioxide 
into  concentrated  solutions  of  sodium  carbonate  : 


NaC0  +  H0 


2  NaHCO 


It  is  an  essential  constituent  of  all  baking  powders. 

Sodium  nitrate  (NaN03).  This  substance,  known  also 
as  Chile  saltpeter,  is  found  in  nature  in  certain  arid 
regions,  where  apparently  it  has  been  formed  by  the 
decay  of  organic  substances  in  the  presence  of  air  and 
sodium  salts.  The  largest  deposits  are  in  Chile,  and  most 
of  the  nitrate  of  commerce  comes  from  that  country. 
Fig.  163  shows  a  deposit  of  sodium  nitrate  in  Chile  after 
it  has  been  broken  apart  by  explosives.  The  commercial 


THE  ALKALI  METALS  405 

salt  is  prepared  by  dissolving  the  crude  nitrate  (known 
as  caliche)  in  water,  allowing  the  insoluble  earthy  mate- 
rials to  settle,  and  evaporating  to  crystallization  the  clear 
solution  so  obtained.  The  soluble  impurities  remain  for 
the  most  part  in  the  mother  liquors. 

Since  this  salt  is  the  only  nitrate  found  extensively  in 
nature,  it  is  the  material  from  which  other  nitrates, 
and  also  nitric  acid,  are  prepared.  Enormous  quantities 
are  used  as  a  fertilizer  and  in  the  manufacture  of  sul- 
furic  acid. 

Sodium  cyanide  (NaCN).  This  salt  of  hydrocyanic  acid 
readily  dissolves  gold  and  is  used  for  extracting  this 
metal  when  it  is  scattered  in  small  quantities  through 
earthy  material.  It  is  prepared  chiefly  by  heating  a  mix- 
ture of  carbon  and  sodamide  (NaNH2),  the  latter  com- 
pound being  obtained  by  the  action  of  sodium  upon 
ammonia  : 

2  Na  +  2  NHS  —  >  2  NaNH2  +  H2 
NaNH+C 


Sodium  cyanide  is  a  white  solid.  Its  aqueous  solution 
is  strongly  alkaline.  The  compound  is  not  only  extremely 
poisonous,  but  in  contact  with  acids  evolves  the  very 
poisonous  hydrocyanic  acid  (HCN). 

The  phosphates  of  sodium.  Sodium  forms  three  salts 
with  phosphoric  acid,  one  normal  and  two  acid  salts. 
The  most  common  of  these  is  disodium  phosphate, 
Na2HPO4.  It  is  prepared  by  the  action  of  phosphoric 
acid  upon  sodium  carbonate  : 

Na2C03  +  H3P04  —  ^Na2HP04  +  H2O  +  CO2 

This  salt  crystallizes  as  the  hydrate,  Na2HPO4-  12H2O. 
It  is  a  constituent  of  the  human  organism. 


406    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  normal  phosphate,  Na3PO4,  is  readily  soluble  in 
water  and  forms  a  strongly  alkaline  solution  due  to  partial 
hydrolysis  : 

NaP0  +  HO  --»-  NaHP0  +  NaOH 


Sodium  hypochlorite  (NaCIO).  This  salt  is  prepared  by 
electrolyzing  a  solution  of  sodium  chloride  under  such 
conditions  that  the  chlorine  and  sodium  hydroxide  gener- 
ated in  the  process  of  electrolysis  react  to  form  sodium 
hypochlorite  as  explained  on  page  276.  The  equation  for 
the  reaction  is  as  follows: 

2  NaOH  +  C12  -  >•  NaCIO  +  NaCl  +  H2O 

The  salt  is  unstable  and  is  obtained  only  in  dilute  solu- 
tion. This  solution  (called  the  Carrel-Dakin  solution)  is 
used  as  an  antiseptic  in  the  treatment  of  wounds. 

Sodium  chlorate  (NaC103).  This  salt  of  chloric  acid  is 
formed  according  to  the  general  method  for  preparing 
chlorates  (p.  276).  When  heated  it  breaks  down  into 
sodium  chloride  and  oxygen,  resembling  potassium  chlo- 
rate in  this  respect  (p.  25).  It  is  an  excellent  oxidizing 
agent  and  is  used  in  the  manufacture  of  fireworks  and 
munitions. 

POTASSIUM 

Occurrence.  Potassium  is  a  rather  abundant  element, 
being  a  constituent  of  many  igneous  rocks,  such  as  the 
feldspars  and  micas.  Very  large  deposits  of  the  chloride 
and  the  sulfate,  associated  with  compounds  of  calcium 
and  magnesium,  are  found  at  Stassfurt,  Germany,  and  are 
known  as  the  Stassfurt  salts. 

The  natural  decomposition  of'  rocks  containing  potas- 
sium gives  rise  to  various  compounds  of  the  element  in  all 


THE  ALKALI  METALS 


407 


fertile  soils.  Its  soluble  compounds  are  absorbed  by  grow- 
ing plants  and  built  up  into  complex  vegetable  tissues ; 
when  these  are  burned,  the  potassium  remains  in  the  ash 
in  the  form  of  carbonate.  The  crude  carbonate  can  be 
separated  from  wood  ashes  by  dissolving  it  in  water.  This 
was  formerly  the  chief  source  of  potassium  compounds, 

but  they  are  now  pre- 

pared  mostly  from  the 
salts  of  the  Stassfurt 
deposits. 

Stas§furt  salts.  These 
salts,  evidently  deposited 
from  sea  water  under 
peculiar  geological  con- 
ditions, form  very  exten- 
sive deposits  in  middle 
and  north  Germany,  the 
most  noted  locality  for 
working  them  being  at 
Stassfurt.  The  deposits 
are  very  thick  and  rest 
upon  an  enormous  layer 
of  common  salt.  They  are 
in  the  form  of  a  series 
of  strata,  each  consisting 

largely  of  a  single  mineral  salt.  Over  thirty  different  minerals 
are  present,  although  some  in  very  small  quantities.  Fig.  164 
shows  a  cross  section  of  these  deposits.  While  from  a  chem- 
ical standpoint  these  strata  are  salts,  they  are  as  solid  and 
hard  as  many  kinds  of  stone  and  are  mined  as  stone  or  coal 
would  be.  Since  the  strata  differ  in  general  appearance,  each 
can  be  mined  separately,  and  the  various  minerals  can  be 
worked  up  by  methods  adapted  to  each  particular  case.  The 
chief  minerals  of  commercial  importance  in  these  deposits 
are  the  following : 


FIG.  164.    Cross-section  diagram  of  the 
Stassfurt  salt  deposits 


408    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Sylvite KC1 

Anhydrite       . CaSO4 

Carnallite KC1  •  MgCl2  •  6  H2O 

Kainite MgSO4  •  KC1  •  3  H2O 

Kieserite    .     .     . '  MgSO4  •  H2O 

Schonite    .     .     '.    ......     K2SO4 -MgSO4  •  6  H2O 

Other  sources  of  potassium.  The  potassium  compounds 
are  of  great  value,  especially  in  the  manufacture  of  ferti- 
lizers. The  importance  of  the  United  States  having  its 
own  source  of  these  compounds  was  made  evident  during 
the  great  war,  when  the  supply  from  the  Stassfurt  salts 
was  cut  off.  The  common  feldspar  which  is  so  abundant 
in  the  earth's  crust  contains  potassium,  but  no  economical 
method  has  been  found  as  yet  for  changing  this  into  the 
desired  compounds,  such  as  potassium  chloride  or  potas- 
sium sulfate.  While  sodium  rather  than  potassium  is  likely 
to  be  present  in  sea  plants,  nevertheless  some  of  these 
plants,  such  as  the  giant  algse  of  the  California  coast, 
contain  potassium  chloride,  amounting  in  some  cases  to 
30  per  cent  of  their,  dry  weight;  indeed,  the  dried  algse 
are  being  used  to  some  extent  as  a  fertilizer.  An  effort 
has  also  been  made  to  obtain  the  salts  from  the  complex 
mineral  alunite  (K2SO4 .  A12(SO4)3 .  4  A1(OH)3),  consid- 
erable deposits  of  which  are  found  in  Utah.  No  satisfac- 
tory source,  however,  has  as  yet  been  found  that  is  at 
all  comparable  with  the  Stassfurt  deposits. 

Preparation  and  properties.  Potassium  is  prepared  by 
methods  similar  to  those  used  in  the  preparation  of  sodium. 
It  is  more  active  than  sodium ;  otherwise  the  properties 
of  the  two  metals  are  alike. 

Potassium  hydroxide  (caustic  potash)  (KOH).  Potas- 
sium hydroxide  is  prepared  by  methods  exactly  similar  to 
those  used  in  the  preparation  of  sodium  hydroxide,  which 


THE  ALKALI  METALS  409 

compound  it  closely  resembles  in  both  physical  and  chemi- 
cal properties.  It  is  not  used  to  any  very  great  extent, 
being  replaced  by  the  cheaper  sodium  hydroxide. 

Action  of  the  halogen  elements  on  bases.  We  have  seen 
that  when  chlorine  is  passed  into  a  solution  of  potassium 
hydroxide  a  reaction  takes  place  (p.  276)  and  that  the 
nature  of  the  reaction  varies  according  to  the  temperature 
of  the  solution.  If  the  solution  is  cold,  potassium  hypo- 
chlorite  and  potassium  chloride  are  formed,  according  to 
the  following  equation : 

2  KOH  +  C12  — >-  KC1O  +  KC1  +  H2O 

If  the  solution  is  hot,  however,  potassium  chlorate  and 
potassium  chloride  are  formed : 

6  KOH  +  3  C12  — >•  KC1O3  +  5  KC1  +  3  H2O 

This  reaction  is  a  general  one  between  the  halogen  ele- 
ments and  the  soluble  bases.  Thus,  in  place  of  chlorine 
in  the  above  equations,  one  may  substitute  bromine  or 
iodine ;  also  in  place  of  the  potassium  hydroxide  one  may 
substitute  sodium  hydroxide  or  calcium  hydroxide.  It  is 
possible  by  this  reaction  to  prepare  a  number  of  important 
compounds.  It  does  not  follow,  however,  that  this  method 
of  preparation  of  any  particular  compound  is  necessarily 
the  most  economical  one. 

Potassium  halides.  Of  these  compounds  potassium  chlo- 
ride is  the  most  familiar,  since  it  is  found  in  such  large 
quantities  in  the  Stassfurt  deposits.  The  mineral  sylvite 
is  nearly  pure  potassium  chloride.  The  salt  is  obtained  not 
only  from  sylvite  but  also  from  carnallite.  In  its  general 
properties  potassium  chloride  resembles  sodium  chloride.  It 
is  used  in  the  preparation  of  nearly  all  other  potassium  salts 
and  as  a  fertilizer.  Potassium  bromide  (KBr)  is  prepared  by 


410    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

the  action  of  bromine  upon  a  hot  solution  of  potassium 
hydroxide  (see  preceding  paragraph).  Potassium  iodide 
(KI)  is  prepared  by  the  same  methods  as  those  used  in 
preparing  potassium  bromide,  with  the  substitution  of 
iodine  for  bromine.  Both  the  iodide  and  the  bromide  are 
used  in  photography  and  in  medicine. 

Potassium  chlorate  (KC103).  This  compound  is  similar 
in  properties  to  sodium  chlorate  (p.  406)  and  is  prepared 
by  the  same  general  methods.  It  is  an  excellent  oxidizing 
agent  and,  like  sodium  chlorate,  is  used  in  the  manufacture 
of  munitions,  fireworks,  and  matches  ;  indeed,  it  is  preferred 
to  sodium  chlorate,  since  the  latter  compound  tends  to 
absorb  moisture  from  the  air. 

Potassium  nitrate  (saltpeter)  (KN03).  This  salt  is  found 
native  in  some  regions  where  the  climate  is  hot  and  dry, 
being  formed  by  the  decay  of  nitrogenous  organic  matter 
in  the  presence  of  earthy  material  containing  potassium. 
Saltpeter  was  formerly  made  by  imitating  these  conditions. 
It  is  now  prepared  by  the  action  of  sodium  nitrate  upon 
potassium  chloride  (the  former  compound  being  obtained 
from  Chile  and  the  latter  from  the  Stassfurt  deposits) : 

NaNO3  +  KC1  — >•  KNO3  +  NaCl 

The  reaction  depends  for  its  success  upon  the  apparently 
insignificant  fact  that  sodium  chloride  is  almost  equally  soluble 
in  cold  and  in  hot  water.  All  four  compounds  represented  in 
the  equation  are  rather  soluble  in  cold  water,  but  in  hot  water 
sodium  chloride  is  far  less  soluble  than  the  other  three.  When 
hot  saturated  solutions  of  sodium  nitrate  and  potassium  chlo- 
ride are  brought  together  sodium  chloride  precipitates  and  can 
be  filtered  off,  leaving  potassium  nitrate  in  solution  together 
with  some  sodium  chloride.  When  the  solution  is  cooled, 
potassium  nitrate  crystallizes  out,  leaving  small  amounts  of 
the  other  salts  in  solution. 


THE  ALKALI  METALS  411 

Potassium  nitrate  is  an  excellent  oxidizing  agent,  and 
its  chief  use  is  in  the  manufacture  of  gunpowder.  For  this 
purpose  it  is  preferable  to  sodium  nitrate,  since  the  latter 
tends  to  absorb  moisture,  and  powder  made  from  it,  if 
exposed  to  air,  soon  becomes  moist  and  unfit  for  use. 
Small  amounts  of  the  nitrate  are  also  used  in  medicine 
and  as  a  preservative  for  meats,  especially  for  corned  beef. 

Potassium  carbonate  (K2C03).  This  compound  can  be 
prepared  from  potassium  chloride  by  the  Leblanc  process, 
just  as  sodium  carbonate  is  prepared  from  sodium  chloride. 
Commercially  it  is  prepared  chiefly  according  to  the 
reactions  indicated  in  the  following  equations : 

3  MgCO,  +  2  KC1  +  C02  +  HO 
^-* 
2  MgKH(CO8)2  — *•  2  MgCO3  +  K2CO3  +  CO2  +  H2O 

Potassium  carbonate  is  used  in  the  manufacture  of  glass. 
When  carbon  dioxide  is  passed  into  an  aqueous  solution 
of  potassium  carbonate,  potassium  bicarbonate  (KHCOg) 
is  formed. 

Other  salts  of  potassium.  Among  the  other  salts  of  potas- 
sium frequently  met  with  are  the  sulfate,  K2SO4 ;  the  acid 
sulfate,  KHSO4 ;  the  acid  sulfite,  KHSO8 ;  and  the  cyan- 
ide, KCN.  They  are  all  white  solids  and  closely  resemble 
the  corresponding  sodium  compounds. 

Deliquescence.  A  large  number  of  salts  resemble  sodium 
nitrate  in  their  tendency  to  absorb  moisture  from  the  air. 
If  the  process  continues  long  enough,  the  salt  may  pass 
into  solution  in  the  water  absorbed.  Such  a  salt  is  said  to 
be  deliquescent,  and  all  very  soluble  salts  are  deliquescent. 
Common  salt  is  not  itself  deliquescent,  but  it  usually  con- 
tarns  enough  deliquescent  salts  of  calcium  and  magnesium 
to  cause  it  to  become  moist  in  wet  weather. 


412    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 
LITHIUM,  RUBIDIUM,  CESIUM 

Of  the  three  remaining  elements  of  the  family  —  lithium, 
rubidium,  and  caesium  —  lithium  is  by  far  the  most  com- 
mon, the  other  two  being  very  rare.  Lithium  chloride  and 
lithium  carbonate  are  not  infrequently  found  in  natural 
mineral  waters,  and  as  these  substances  are  supposed  to 
increase  the  medicinal  value  of  the  water,  small  quantities 
of  them  are  often  added  to  artificial  mineral  waters. 

COMPOUNDS  OF  AMMONIUM 

General.  As  explained  in  Chapter  XVIII,  when  ammo- 
nia is  passed  into  water  the  two  unite  to  form  the  base 
ammonium  hydroxide,  and  when  this  base  is  neutralized 
with  acids,  ammonium  salts  are  formed.  Since  the  ammo- 
nium group  is  univalent,  ammonium  salts  resemble  those 
of  the  alkali  metals  in  formulas ;  they  also  resemble  the 
latter  salts  in  their  chemical  properties,  and  may  be  con- 
veniently described  in  connection  with  them.  They  all 
volatilize  upon  being  heated,  most  of  them  being  decom- 
posed in  the  process.  When  heated  with  an  aqueous 
solution  of  sodium  hydroxide,  they  evolve  ammonia 
(p.  201).  Since  the  ammonia  can  be  easily  recognized, 
the  reaction  serves  for  the  detection  of  the  presence  of 
ammonium  compounds. 

Occurrence.  Small  quantities  of  ammonium  compounds 
are  found  in  the  soil.  They  are  being  continually  absorbed 
by  growing  plants,  but  are  returned  to  it  again  in  the 
process  of  decay.  They  are  also  found  in  sea  water  and 
in  some  volcanic  regions.  Larger  quantities  are  found  in 
the  Stassfurt  deposits.  Commercially  ammonium  com- 
pounds are  prepared  from  the  ammoniacal  liquors  pro- 
duced in  the  manufacture  of  coke  (pp.  306,  307). 


THE  ALKALI  METALS  413 

Ammonium  chloride  (sal  ammoniac)  (NH4C1).  This  is  a 
white  solid.  When  heated  it  partly  decomposes  into  ammonia 
and  hydrogen  chloride,  which  recombine,  as  the  temperature 


This  salt  is  used  in  soldering,  since  the  hydrogen  chloride 
evolved  by  the  heat  removes  the  oxide  from  the  surface 
of  the  metals.  It  is  also  used  in  making  dry  cells,  in 
medicine,  and  as  a  chemical  reagent. 

Ammonium  sulfate,  (NH4)2S04.  This  salt  resembles  the 
chloride  very  closely  and,  being  cheaper,  is  used  in  place 
of  it  when  possible.  It  is  used  in  large  quantities  as  a 
fertilizer,  the  nitrogen  which  it  contains  being  a  very 
valuable  food  for  plants. 

The  carbonates  of  ammonium.  Both  the  normal  carbonate, 
(NH4)2CO3,  and  the  acid  carbonate,  NH4HCO3,  are  white 
solids,  readily  soluble  in  water.  The  normal  carbonate  slowly 
decomposes  into  the  acid  carbonate,  evolving  ammonia  : 

(NH4)2C03  —  >-  NH4HCO3  +  NH3 

Ammonium  sulfides.  When  hydrogen  sulfide  is  passed 
into  aqua  ammonia  a  solution  containing  ammonium  sul- 
fide, (NH4)2S  and  ammonium  acid  sulfide  (NH4HS)  is 
obtained  : 

NH4OH  +  H2S  —  »-NH4HS  +  H2O 
2  NH4OH  +  H2S  —  >-  (NH4)2S  +  2  H2O 

This  solution  is  usually  known  simply  as  ammonium  sulfide, 
and  is  used  as  a  reagent  in  testing  for  certain  metals.  When 
exposed  to  the  air  it  slowly  decomposes,  and  the  sulfur  lib- 
erated in  the  process  combines  with  the  compounds  present, 
forming  different  sulfides,  such  as  (NH^S,,  and  (NH4)2S8  or, 
in  general,  (NH4)2S.,..  The  resulting  solution  is  yellow  and  is 
termed  yellow  ammonium  suljide  or  ammonium  polysulfide. 


414    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Flame  reactions.  There  are  several  metals  which,  when 
volatilized  in  a  colorless  flame,  such  as  that  of  a  Bunsen 
burner,  impart  a  characteristic  color  to  the  flame.  Thus, 
sodium  (or  any  of  its  compounds  that  will  volatilize  in 
the  heat  of  the  flame)  imparts  to  the  flame  a  strong  yellow 

color.  Potassium  and 
its  compounds  color  the 
flame  a  pale  violet,  and 
lithium  colors  it  a  deep 
crimson  red. 

Advantage  is  taken 
of  these  facts  in  testing 
for  the  presence  of  the 
elements  in  different 
substances.  The  test  is 
FIG.  165.  Method  of  making  a  flame  test  best  made  by  using  a 

platinum  wire  one  end 

of  which  is  fused  into  a  piece  of  glass  tubing  that  serves 
as  a  handle.  The  other  end  of  the  wire  is  dipped  into  water 
and  rubbed  in  the  substance  to  be  tested  (or  dipped  into 
a  concentrated  solution  of  the  substance),  and  the  wire 
with  the  adhering  particles  is  Field  in  the  outer  edge  of 
the  base  of  the  Bunsen  flame  (Fig.  165). 


1.  What  is  an  alkali?   Can  a  metal  itself  be  an  alkali? 

2.  Write    equations    showing  how  the    following  changes  may 
be  brought  about,  giving  the  general  principle  involved  in  each 
change:   NaCl  -  >-Na2SO3,  Na2SO3  -  >-NaCl,  NaCl 


Na2SO4 


NaNO 


NaHCO 


3.  What  carbonates  are  soluble  ? 

4.  State  the   conditions  under  which  the  reaction  represented. 
by  the  following  equation  can  be  made  to  go  in  either  direction  : 

Na2CO8  +  H2O  +  CO2^=>:2  NaHCO8 


THE  ALKALI  METALS  415 

5.  Account  for  the  fact  that  solutions  of  sodium  carbonate  and 
potassium  carbonate  are  alkaline. 

6.  What  nonmetallic  element  is  obtained  from  the  deposits  of 
Chile  saltpeter  ? 

7.  Supposing  concentrated  hydrochloric  acid  (den.  =  1.2)  to  be 
worth  eight  cents  a  pound,  what  is  the  value  of  the  acid  generated  in 
the  preparation  of  1  ton  of  sodium  carbonate  by  the  Leblanc  process  ? 

8.  What  weight  of   sal   soda  can  be  prepared  from  1  kg.  of 
anhydrous  sodium  carbonate? 

9.  Write  equations  for  the  preparation  of  potassium  hydroxide 
by  three  different  methods. 

10.  What  would  take  place  if  a  bit  of  potassium  hydroxide  were 
left  exposed  to  the  air  ? 

11.  Write  the  equations  for  the  reactions  between  sodium  hydrox- 
ide and  bromine ;  between  potassium  hydroxide  and  chlorine. 

12.  Write  equations  for  the  preparation  of  potassium  sulf ate ; 
of  potassium-acid  carbonate. 

13.  What  weight  of  carnallite  would  be  necessary  in  the  prepara- 
tion of  1  ton  of  potassium  carbonate  ? 

14.  AVrite  the  equations  showing  how  ammonium  chloride,  ammo- 
nium sulfate,  ammonium  carbonate,  and  ammonium  nitrate  may  be 
prepared  from  ammonium  hydroxide. 

15.  Write  an  equation  to  represent  the  reaction  involved  in  the 
preparation  of  ammonia  from  ammonium  chloride. 

16.  What    substances    already   studied    are   prepared   from    the 
following  compounds :    ammonium   chloride ;    ammonium   nitrate  ; 
ammonium  nitrite ;    sodium  nitrate ;    sodium  chloride  ? 

17.  Write  equations  for  the  preparation  of  potassium  iodide; 
potassium  bromide. 

18.  How  could  you  distinguish  between  potassium  chloride  and 
potassium  iodide?  between  sodium  chloride  and  ammonium  chloride? 
between  sodium  nitrate  and  potassium  nitrate  ? 

19.  What  are  the  relative  advantages  of  sodium  chlorate  and 
potassium  chlorate  as  oxidizing  agents? 

20.  Write  the  names  and  formulas  for  the  different  compounds 
that  may  be  formed  by  the  action  of  the  halogen  elements  on  the 
soluble  bases. 


CHAPTER  XXXIII 
SOAP;  GLYCERIN;  EXPLOSIVES 

Introductory.  At  first  thought  one  might  wonder  why 
three  products  so  different  from  each  other  as  are  soap, 
glycerin,  and  explosives  should  be  brought  together  for 
study  in  the  same  chapter.  However,  the  grouping  is  a 
natural  one  industrially,  for  glycerin  is  a  by-product  in 
the  manufacture  of  soap;  and  nitroglycerin,  one  of  the 
most  powerful  explosives,  is  prepared  from  glycerin  and 
stands  in  a  general  way  as  a  type  of  an  explosive  com- 
pound. It  is  convenient,  therefore,  in  a  very  brief  de- 
scription of  these  products,  to  include  all  three  in  the 
same  chapter. 

Composition  of  soap,  and  materials  used  in  its  preparation. 
Soap  is  a  mixture  of  the  sodium  and  potassium  salts  of 
oleic,  palmitic,  and  stearic  acids  (p.  342).  The  essential 
materials  used  in  the  preparation  of  soap  are  as  follows: 

1.  Fat  or  oil.    As  shown  on  page  342,  fats  and  oils  are 
largely    mixtures    of    olein,   palmitin,    and    stearin.    The 
cheaper  grades  of  these  are  used  in  making  soap.    Those 
commonly  employed  are  a  low  grade  of  animal  fat  (tallow) 
and  the   cheaper  vegetable  oils,  such  as   cottonseed  oil, 
coconut  oil,  and  palm  oil. 

2.  Alkali.    The  alkali  used  is  the  hydroxide  of  either 
sodium  or  potassium.    Sodium  hydroxide  is  nearly  always 
used,  since  it  gives  a  hard  soap,  while  potassium  hydroxide 
gives  a  soft  soap.    Other  bases,  such  as  calcium  hydroxide, 
are  unavailable  since  they  yield  insoluble  soaps. 

416 


SOAP;   GLYCERIN;   EXPLOSIVES 


417 


Reaction  taking  place  in  the  preparation  of  soap.  When 
the  fat  and  the  alkali  are  heated  together  under  proper 
conditions  the  olein,  palmitin,  and  stearin  present  in  the 
fat  are  decomposed,  forming  glycerin  together  with  sodium 
oleate,  sodium  palmitate,  and  sodium  stearate.  A  mixture 
of  these  three  salts  constitutes  soap.  The  reactions  may 
be  illustrated  by  the  following  equation,  which  represents 
the  change  taking  place  when  stearin  is  heated  with 
sodium  hydroxide: 


C8H5(C18H3602)3+3  NaOH  —  ^ 

stearin  sodium  hydroxide  glycerin  sodium  stearate 

In  this  reaction  the  fat  is  said  to  be  saponified,  and  the 
process  is  known  as  saponification. 

Commercial  manufacture  of  soap. 
The  oil  or  melted  fat  is  poured  into 
large  iron  kettles  together  with  a 
solution  of  sodium  hydroxide  con- 
taining about  one  fourth  of  the 
amount  of  alkali  necessary  to  sa- 
ponify the  fat.  As  a  rule  the  kettles 
(Fig.166)  are  very  large,  500,000  11>. 
or  more  of  soap  being  made  in  some 
of  them  in  a  single  heating.  They 
are  provided  with  coils  of  steam 
pipe  for  heating  the  mixture.  The 
fat  and  alkali  are  stirred  by  forcing 
air  or  live  steam  into  the  bottom 
of  the  mixture.  As  the  heating 
continues,  the  remainder  of  the 
alkali  is  added.  The  reaction  is 
complete  in  from  two  to  five  days. 

The  soap  is  next  removed,  or  salted  out,  from  the  mixture. 
This  process  consists  in  adding  salt  and  again  heating.  After 
a  time  the  soap  rises  to  the  top  of  the  liquid,  or  spent  lye,  as 
it  is  called.  The  soap  so  obtained  is  purified  and  then  run 


FIG.  166.    A  soap  kettle 


418    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

into  a  mixing  machine  (crutcher).  Here  it  is  mixed  with 
any  appropriate  material  which  it  is  desired  to  add,  such 
as  perfume,  borax,  sodium  silicate,  or  sodium  carbonate.  It 
is  then  run  into  large  molds  to  harden,  after  which  it  is 
cut  and  pressed  into  cakes  of  the  desired  size.  The  glycerin 
formed  in  the  reaction  is  separated  from  the  spent  lye  by 
distillation. 

Varieties  of  soap.  Transparent  soaps  are  ordinarily  made  by 
dissolving  soap  in  alcohol.  The  solution  is  filtered  and  the 
excess  of  alcohol  removed  by  distillation.  Castile  soaps  are 
made  from  mixtures  of  olive  oil  with  cheaper  oils.  The  color 
of  mottled  soaps  is  produced  by  the  addition  of  ferrous  sulfate, 
Prussian  blue,  or  some  similar  pigment.  Floating  soaps  owe 
their  lightness  to  bubbles  of  air.  Naphtha  soaps  contain  from 
5  per  cent  to  10  per  cent  of  petroleum  naphtha.  Scouring 
soaps  contain  from  5  per  cent  to  10  per  cent  of  soap  and 
from  80  per  cent  to  90  per  cent  of  some  abrasive  material  such 
as  fine  sand  or  volcanic  ash.  Sometimes  a  small  percentage 
of  sodium  carbonate  is  also  present.  Soap  poivders  are,  as  a 
rule,  sodium  carbonate  or  a  mixture  of  sodium  carbonate  and 
ground  soap. 

Properties  of  soap.  Soap  dissolves  in  soft  waters,  giving 
a  slightly  alkaline  solution  due  to  hydrolysis.  If  an  acid, 
such  as  hydrochloric  acid,  is  added  to  the  aqueous  solu- 
tion, the  organic  acids  are  liberated  from  their  salts  and 
are  precipitated  in  the  form  of  white  insoluble  solids: 

NaC^O,  +  HC1 >•  NaCl  +  H  •  C  HO2 

sodium  stearate  stearic  acid 

The  calcium  and  magnesium  salts  of  oleic,  palmitic,  and 
stearic  acids  are  insoluble  in  water  and  are  therefore  pre- 
cipitated when  a  calcium  or  magnesium  compound  is  added 
to  an  aqueous  solution  of  soap  : 

2  NaCMH  O  +  CaCl2 *  2  NaCl  +  Ca(ClgH;i6O2)2 

sodium  stearate  "  calcium  stearate 


SOAP;  GLYCEKIN;  EXPLOSIVES  419 

It  is  due  to  this  fact  that  soaps  do  not  lather  with  hard 
waters  (p.  431),  but  form  a  curdy  precipitate,  since  such 
waters  always  contain  salts  of  calcium  and  magnesium  in 
solution. 

Cleansing  action  of  soap.  Attention  has  been  called  to 
the  property  possessed  by  soap  of  aiding  in  the  formation 
of  emulsions  (p.  387).  The  cleansing  power  of  soap  is 
largely  due  to  this  property.  When  soap  is  rubbed  upon 
the  skin  any  oily  substances  present  are  emulsified  by  the 
soap  and  washed  away. 

Glycerin,  C3H6(OH)3.  This  is  a  colorless  oily  liquid  hav- 
ing a  sweet  taste.  It  is  formed  whenever  a  fat  is  acted 
upon  by  an  alkali  and  consequently  is  a  by-product  in  the 
manufacture  of  soap.  Nitric  acid  acts  upon  it,  forming 
glyceryl  nitrate,  as  indicated  in  the  following  equation : 

C8H5(OH)3  +  3  HN03  — >-  C8H5(N03)8  +  3  H2O 

In  actually  carrying  out  this  reaction  a  mixture  of  nitric , 
and  sulfuric  acid  is  always  used ;  the  sulfuric  acid  aids  in 
the  reaction  by  absorbing  the  water  produced.  It  will  be 
noted  that  the  reaction  is  exactly  similar  to  that  of  nitric 
acid  upon  a  base,  so  that  glyceryl  nitrate  is  really  a  salt 
of  nitric  acid.  Glyceryl  nitrate  is  a  slightly  yellowish  oil. 
It  is  a  powerful  explosive  and  is  the  chief  constituent  of 
the  explosive  known  as  nitroglycerui.  The  chief  use  of 
glycerin  is  in  the  preparation  of  this  product. 

Explosives.  An  explosion  is  caused  by  a  very  rapid 
chemical  reaction  which  results  in  the  formation  of  large 
volumes  of  gases  from  liquid  and  solid  materials  called 
explosives.  The  greater  the  volume  change  and  the  more 
rapidly  it  is  produced,  the  more  violent  the  explosion. 
The  most  common  of  the  manufactured  explosives  are 
the  following: 


420    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

1.  Black  gunpowder.    Ordinary  black  gunpowder  is  an  inti- 
mate mixture  of  potassium  nitrate,  sulfur,  and  charcoal.    When 
this  mixture  is  ignited  complicated  reactions  occur,  the  prin- 
cipal change  being  indicated  in  the  following  equation : 

2  KN03  +  3  C  +  S >-K?S  +  3  CO2  +  N2 

The  explosive  effects  are  due  to  the  sudden  liberation  of 
highly  heated  gaseous  products,  especially  carbon  dioxide  and 
nitrogen.  Less  than  50  per  cent  by  weight  of  the  products 
formed  in  the  explosion  are  gases,  so  that  black  gunpowder  is 
not  comparatively  a  powerful  explosive. 

2.  Nitroglycerin.    As  stated  above,  nitroglycerin  is  made  by 
the  action  of  a  mixture  of  nitric  and  sulfuric  acids  upon  glyc- 
erin and  is  composed  chiefly  of  glyceryl  nitrate.    It  is  one  of 
our  most  powerful  explosives  —  much  more  powerful  than  black 
gunpowder.    Tlie  changes  that  take  place  in  its  decomposition 
are  represented  in  a  general  way  in  the  following  equation : 

4  C3H5(N03)3 » 12  C02  +  6  N2  +  10  H20  +  02 

-One  volume  of  nitroglycerin  yields  on  explosion  about  1300 
volumes  of  gas,  which  is  expanded  by  the  heat  of  the  reaction 
to  over  10,000  volumes.  Pure  nitroglycerin  is  very  dangerous 
because  of  the  ease  with  which  it  is  set  off.  Large  quantities 
are  used  in  making  dynamite,  in  which  form  it  is  not  exploded 
so  readily  by  jarring  and  can  be  transported  with  less  danger. 
Ordinary  dynamite  consists  of  a  mixture  of  sodium  nitrate, 
nitroglycerin,  and  wood  pulp,  the  latter  acting  as  an  absorbent 
for  the  nitroglycerin. 

3.  Nitrocellulose.    Just  as  glycerin  is  changed,  into  nitro- 
glycerin by  the  action  of  a  mixture  of  nitric  and  sulfuric  acids, 
so  cellulose  (p.  328)  by  a  similar  treatment  yields  nitrocellulose. 
Nitrocellulose  is  prepared  from  a  pure  form  of  cellulose,  such 
as  cotton,  and  resembles  cotton  in  appearance.   Like  nitroglyc- 
erin, it  is  a  far  more  powerful  explosive  than  black  gunpowder. 
If  ignited  under  ordinary  conditions,  it  will  burn  quietly.    If 
subjected  to  a  sudden  shock  (such  as  may  be  produced  by  the 
explosion   of  a  small   percussion   primer)  the   nitrocellulose 


SOAP;  GLYCERIN;  EXPLOSIVES  421 

decomposes  with  enormous  violence.  The  products  of  the 
decomposition  are  all  colorless  gases ;  hence  the  use  of  this 
explosive  in  making  smokeless  gunpowder.  When  nitrocellu- 
lose is  used  for  this  purpose  it  is  necessary  to  modify  the 
pure  material  somewhat,  as  otherwise  the  violence  of  the  ex- 
plosion would  shatter  any  firearms  in  which  the  powder  was 
used.  This  is  done  by  mixing  nitrocellulose  with  sufficient 
alcohol  and  ether  to  form  a  plastic  mass,  which  is  then  molded 


Fi<i.  167.    Powder  grains  for  large  guns  (natural  size) 

into  the  form  of  rods  (grains)  with  a  number  of  perforations 
through  the  rods.  Fig.  167  shows  the  form  of  the  grains  used 
in  the  large  guns  of  our  navy.  Gelatin  dynamite  is  made  by 
stirring  nitrocellulose  into  nitroglycerin.  It  forms  a  jellylike 
mass  and  is  a  very  powerful  explosive. 

4.  Trinitrotoluene,  C7H5(N02)3 ;  picric  acid,  C6H2(N02)3OH.  In 
connection  with  the  great  war  there  has  been  developed  a  power- 
ful explosive  known  as  trinitrotoluene.  It  is  prepared  by  the 
action  of  nitric  acid  on  toluene  (p.  336).  It  is  a  white  solid  and 
can  be  transported  with  safety.  Picric  acid  is  also  a  powerful 
explosive.  It  is  a  yellow  solid  made  by  treating  phenol  (p.  337) 
with  nitric  acid. 

EXERCISES 

1.  In  what  way  does  aqua  ammonia  assist  in  the  removal  of  grease  ? 

2.  For  what  is  lye  used  as  a  household  article  ? 

3.  What  effect  will  the  softening  of  a  city  water  supply  have  on 
soap  consumption? 


422    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

4.  Why  will  gas  which  burns  quietly  in  a  stove  explode  violently 
if  a  sufficient  quantity  of  it  is  allowed  to  escape  into  a  room  and  is 
then  ignited  ? 

5.  What  is  the  significance  of  the  word  glycerin'! 

6.  Why  not  use  sodium  nitrate  in  making  gunpowder? 

7.  What  are  the  approximate  proportions  in  which  the  constitu- 
ents of  gunpowder  are  mixed  in  its  preparation  (see  equation  for 
reaction  of  explosion)  ? 

8.  Why  are  some  gunpowders  smokeless  and  others  not? 

9.  In  some  cities  garbage  is  being  utilized  as  a  source  of  glyc- 
erin.   Suggest  the  chemistry  involved  in  the  process. 

10.  Why  not  use  dynamite  as  an  explosive  in  guns  ? 

11.  Why  does  the  removal  of  water  by  sulfuric  acid  assist  in 
making  nitroglycerin  ? 


CHAPTER  XXXIV 
THE  CALCIUM  FAMILY;    FERTILIZERS 


NAME 

SYMBOL 

ATOMIC   ni,VSTTV 

MELTING 

CARBONATE 

WEIGHT 

POINT 

DECOMPOSES 

Calcium       .     .     . 

Ca 

40.07 

1.55 

810° 

at  dull-red  heat 

Strontium    .     .     . 

Sr 

87.63 

2.54 

at  white  heat 

Barium  .... 

Ba 

137.37 

3.75 

850° 

scarcely  at  all 

The  family.  The  calcium  family  consists  of  the  very 
abundant  metal  calcium  and  the  rarer  metals  strontium 
and  barium.  These  three  metals  are  often  termed  the  alka- 
line earth  metals.  Radium  also  occurs  in  this  family  (see 
the  periodic  arrangement  of  the  elements),  but  it  is  more 
convenient  to  discuss  it  in  connection  with  the  element 
uranium,  to  which  it  bears  a  peculiar  relation. 

1.  Occurrence.    Like  the  alkali  metals,  the  alkaline  earth 
metals  do  not  occur  free  in  nature.    Their  most  abundant 
compounds  are  the  carbonates  and  sulfates,  calcium  also 
occurring  in   large   quantities   in   the   form  of  the  phos- 
phates and  silicates. 

2.  Preparation.     The  metals  are  prepared  by  the  elec- 
trolysis of  their  melted  chlorides  or  hydroxides.    Calcium 
is  the  most  easily  prepared. 

3.  Properties.     The  three  metals  resemble  one  another 
very   closely.    They  are   silvery  white  in   color  and   are" 
somewhat  harder  than  lead.    They  decompose  water  at 
ordinary  temperatures,  forming  hydroxides  and  liberating 
hydrogen,  although  not  so  readily  as  do  the  alkali  metals. 

423 


424    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

When  ignited  in  air  they  burn  with  brilliancy,  forming  ox- 
ides. These  oxides  combine  with  water  to  form  hydroxides. 
4.  Compounds.  The  alkaline  earth  metals  act  as  bivalent 
elements  in  the  formation  of  salts.  The  corresponding 
salts  of  the  three  metals  are  similar  to  one  another  and 
show  a  regular  gradation  in  many  of  their  properties. 
Unlike  the  alkali  metals,  their  carbonates,  phosphates, 
and  silicates  are  insoluble  in  water.  Barium  sulfate  is 
also  insoluble,  while  the  sulfates  of  calcium  and  strontium 
are  but  sparingly  soluble.  When  volatilized  in  a  colorless 
flame  (p.  414)  the  compounds  of  each  of  the  three  metals 
impart  a  characteristic  color  to  the  flame,  those  of  calcium 
making  the  flame  orange  in  color,  those  of  strontium  mak- 
ing it  crimson,  and  those  of  barium  making  it  green. 

CALCIUM 

Occurrence.  Calcium  is  one  of  the  abundant  elements. 
In  the  form  of  a  carbonate  it  occurs  in  a  number  of 
different  forms,  such  as  limestone  and  marble.  The  most 
important  of  its  mineral  compounds  are  the  following: 

Calcite      ....     CaCO3  Wollastonite    .     CaSiO3 

Phosphorite .     .     .     Ca3(PO4)2          Gypsum  .     .     .     CaSO4  •  2  H2O 
Fluorite   ....     CaF2  Anhydrite    .     .     CaSO4 

Preparation.  Calcium  is  prepared  commercially  by  the 
electrolysis  of  the  melted  chloride  in  the  following  way : 

The  apparatus  consists  of  a  cylindrical  iron  vessel  (Fig.  168), 
through  the  bottom  of  which  extends  the  iron  cathode  A .  The 
anodes  (B,  B),  several  in  number,  are  placed  about  the  sides  of 
the  vessel.  The  calcium  separates  in  a  molten  condition  at  the 
cathode  A  and  rises  in  the  form  of  globules  to  the  lower  surface  of 
a  solid  stick  of  calcium  (Z>),  suspended  above  the  cathode.  There 
it  is  chilled  by  a  water-cooling  device  C,  C,  and  adheres  to  the 
stick  of  calcium,  which  is  slowly  raised  as  it  increases  in  length. 


THE  CALCIUM  FAMILY;  FERTILIZERS       425 


Properties.  Calcium  is  a  silver-white  metal  only  a  little 
heavier  than  water  and  melting  at  810°.  It  combines 
readily  with  many  of  the  elements  and  when  ignited 
burns  in  oxygen  with  dazzling  brilliancy.  Like  sodium 
and  potassium,  it  decomposes  water,  forming  a  hydroxide 
and  hydrogen.  As  yet  it  has  few  commercial  applications. 

Calcium  oxide  (lime)  (CaO). 
Pure  calcium  oxide  can  be 
prepared  by  burning  calcium 
in  oxygen  or  by  heating  the 
nitrate  or  the  carbonate.  The 
more  or  less  impure  oxide, 
known  commercially  as  lime 
or  quicklime,  is  prepared  on 
a  large  scale  by  heating 
limestone,  which  is  chiefly 
calcium  carbonate  (CaCO3). 
When  heated  the  calcium 
carbonate  decomposes  ac- 
cording to  the  following 
equation : 

CaCO3q=»:CaO  +  CO2 

The  reaction  is  reversible,  so  that  in  manufacturing  practice 
the  decomposition  is  effected  under  conditions  that  will 
conduct  away  the  carbon  dioxide  as  fast  as  it  is  formed. 
Ordinary  lime  is  a  white  amorphous  substance.  When 
heated  intensely,  as  in  the  oxyhydrogen  flame,  it  gives  a  bril- 
liant light  called  the  limelight.  It  melts  only  in  the  heat  of 
the  electric  furnace.  Water  acts  upon  lime  with  the  evolu- 
tion of  a  great  deal  of  heat, — whence  the  name  quicklime,  or 
live  lime,  —  the  process  being  called  slaking.  The  equation  is 

CaO  +  H2O  — >•  Ca(OH)2  +  15,540  cal. 


FIG.  168.    The  preparation  of  cal- 
cium by  electrolysis 


426    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Because  of  its  affinity  for  water  it  is  used  for  drying  gases. 
It  also  absorbs  carbon  dioxide,  forming  the  carbonate : 


CaO  +  CO0 


•CaCO0 


Lime  exposed  to  air  is  therefore  gradually  converted  into 
the  hydroxide  and  the  carbonate  and  will  no  longer  slake 
with  water.  It  is  then  said  to  be  air-slaked.  Lime  is  pro- 
duced in  enormous  quantities  for  use 
in  making  calcium  hydroxide. 

Commercial  production  of  lime.  A  vertical 
section  of  the  newer  form  of  limekiln  is 
shown  in  Fig.  169.  The  kiln  is  about 
50  ft.  in  height.  A  number  of  fire  boxes, 
or  furnaces  (A,  AJ,  are  built  around  the 
lower  part,  all  leading  into  the  central 
stack.  The  kiln  is  filled  with  limestone 
through  a  swinging  door  B.  The  hot 
products  of  combustion  are  drawn  up 
through  the  kiln,  and  the  limestone  is 
gradually  decomposed  by  the  heat.  The 
bottom  of  the  furnace  is  so  constructed 
that  a  current  of  air  is  drawn  in  at  C. 
FIG.  169.  The  vertical  This  seryes  the  PurPose  of  cooling  the 
section  of  a  modem  hot  lime  at  the  base  of  the  furnace,  of 
limekiln  furnishing  heated  oxygen  for  the  com- 

bustion,  and    of    removing    the    carbon 

dioxide  from  the  kiln.  The  lime  is  dropped  into  cars  run  under 
the  furnace.  Generally  a  number  of  these  kilns  are  operated 
together,  as  shown  in  Fig.  170. 

Calcium  hydroxide  (slaked  lime),  Ca(OH)a.  This  com- 
pound is  prepared  by  adding  water  to  lime,  as  explained 
above.  When  pure  it  is  a  light  white  powder.  It  is 
sparingly  soluble  in  water.  Moreover,  its  solubility  dimin- 
ishes with  rise  in  temperature.  The  aqueous  solution  is 
called  limewater.  Owing  to  its  cheapness  calcium  hydroxide 


THE  CALCIUM  FAMILY;  FERTILIZERS       427 


is  used  in  the  industries  whenever  an  alkali  is  desired. 
It  is  used  in  the  preparation  of  ammonia,  bleaching  powder, 
and  the  hydroxides  of  sodium  and  potassium.  It  is  also 
used  to  remove  sulfur  compounds  and  carbon  dioxide  from 
coal  gas,  to  remove  the  hair  from  hides  in  making  leather, 
for  making  mor- 
tar and  plaster, 
and  for  liming 
soils  (p.  437). 

Mortar  and  plas- 
ter. Mortar  is  a 
mixture  of  cal- 
cium hydroxide 
and  sand.  When 
it  is  exposed  to 
the  air  or  spread 
upon  porous  ma- 
terials moisture 
is  removed  from 
it  (partly  by  ab- 
sorption into  the 
porous  materials 

and  partly  by  evaporation)  and  the  mortar  becomes  firm,  or  sets. 
At  the  same  time  carbon  dioxide  is  slowly  absorbed  from  the 
air,  and  hard  calcium  carbonate  is  formed  : 


FIG.  170.   A  group  of  limekilns  in  a  modern  plant 


CaCO 


H20 


Ca(OH)2  +  C02 

By  this  combined  action  the  mortar  becomes  very  hard  and 
adheres  firmly  to  the  surface  upon  which  it  is  spread.  The 
sand  serves  to  give  body  to  the  mortar  and  makes  it  porous  ; 
it  also  prevents  too  much  shrinkage.  Plaster  is  a  mixture  of 
calcium  hydroxide  and  hair,  the  latter  serving  to  hold  the 
mass  together. 

Bleaching  powder.    When  chlorine  is  passed  over  calcium 
hydroxide  there  is  formed  a  white  solid  compound  having 


428    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

the  formula  CaOCl2  and  known  as  bleaching  powder,  chlo- 
ride of  lime,  or  simply  bleach : 

Ca(OH)2  +  C12 — )-CaOCl2  +  H2O 

When  this  compound  is  treated  with  an  acid,  chlorine  is 
evolved : 

CaOCl2  +  H2SO4  -3-*  CaSO4  +  H2O  +  Cla 

When  exposed  to  the  air,  bleaching  powder  is  slowly  acted 
upon  by  moisture  and  carbon  dioxide,  with  the  liberation 
of  hypochlorous  acid  (HC1O),  which  is  a  good  disinfectant. 


FIG.  171.   Diagram  of  a  plant  for  making  bleaching  powder 

Bleaching  powder  is  prepared  in  large  quantities  for  use 
as  a  bleaching  agent,  as  a  disinfectant,  and  as  an  agent  for 
purifying  city  water  supplies.  The  commercial  product  gener- 
ally contains  from  35  to  37  per  cent  of  available  chlorine. 

In  the  commercial  preparation  of  bleaching  powder  the  cal- 
cium hydroxide  is  spread  to  a  depth  of  2  or  3  inches  upon  the 
floor  of  a  room,  usually  made  of  lead  or  concrete  (Fig.  171). 
The  chlorine,  generated  by  the  electrolysis  of  sodium  chloride, 
enters  near  the  top  at  A.  Any  unabsorbed  chlorine  passes  out 
at  B  and  into  the  adjoining  chamber  at  f. 

Calcium  carbonate  (CaC03).  Enormous  quantities  of  cal- 
cium carbonate  occur  in  nature.  Limestone  is  the  most 
abundant  form  and  sometimes  constitutes  whole  mountain 
ranges.  Limestone  is  never  pure  calcium  carbonate,  but 
contains  variable  percentages  of  magnesium  carbonate,  clay, 
silica,  and  compounds  of  iron.  Pearls,  coral,  and  various 


THE  CALCIUM  FAMILY;  FERTILIZERS       429 

shells  are  largely  calcium  carbonate.  Calcite  is  a  very 
pure,  crystalline  form  and  often  is  found  in  large  trans- 
parent crystals  (Fig.  172)  called  Iceland  spar.  Marble  is 
composed  of  very  small  calcite  crystals. 

Calcium  carbonate  is  an  example  of  a  dimorphous  compound ; 
that  is,  it  crystallizes  in  two  different  forms.  Thus,  calcite  crystals 
belong  to  the  hexagonal  sys- 
tem (see  Appendix),  while 
aragonite,  another  form  found 
in  nature,  forms  crystals  be- 
longing to  the  rhombic  system. 

In  the  laboratory  pure 
calcium  carbonate  can  be 
prepared  by  treating  a  sol- 
uble calcium  salt  with  a 
soluble  carbonate: 


CaCOa  +  2NaCl 

When  prepared  in  this  way  it  is  a  soft  white  powder  often 
called  precipitated  chalk  and  is  much  used  as  a  polishing 
powder  (tooth  powder). 

The  natural  varieties  of  calcium  carbonate  find  many 
uses,  as  in  the  preparation  of  lime  and  of  carbon  dioxide ; 
in  metallurgical  operations,  especially  in  blast  furnaces ;  in 
the  manufacture  of  soda  and  glass ;  for  building  stone ; 
and  for  ballast  for  roads. 

Calcium  acid  carbonate,  Ca(HC03)a.  Calcium  carbonate  is 
almost  insoluble  in  pure  water.  It  readily  dissolves,  however, 
in  water  which  holds  carbon  dioxide  in  solution.  This  is  due 
to  the  fact  that  the  carbonate  combines  with  the  carbonic  acid 
present  in  the  water  and  forms  calcium  acid  carbonate,  which 

CaCO3  +  H2CO3  ^n 


430    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  resulting  acid  carbonate  exists  only  in  solution,  since 
it  is  unstable  and  decomposes  into  the  normal  carbonate 
on  heating  or  on  evaporation  of  its  solution. 

Natural  waters  always  contain  more  or  less  carbon  dioxide 
in  solution.  In  the  case  of  certain  underground  waters  the 
amount  of  carbon  dioxide  is  comparatively  large,  being  held 
in  solution  by  pressure.  Such  waters  have  a  marked  solvent 
action  upon  limestone,  dissolving  both  the  calcium  carbonate 
and  the  magnesium  carbonate.  In  certain  localities  this  sol- 
vent action,  continued  through  geological  ages,  has  resulted  in 
the  formation  of  large  caves  in  limestone  rock,  such  as  the 
Mammoth  Cave  in  Kentucky. 

Calcium  sulfate  (CaS04).  Calcium  sulfate  occurs  in  nature 
in  several  different  forms,  the  most  common  of  which  is 
gypsum  (CaSO4  •  2  H2O).  This  is  quarried  in  large  amounts 
in  New  York,  Michigan,  and  Oklahoma.  It  is  used  as  a  filler 
in  making  paper  (p.  330),  as  a  constituent  of  fertilizers, 
and  especially  in  making  plaster  of  Paris. 

Calcium  sulfate  is  but  slightly  soluble  in  water  and  is 
precipitated  in  the  form  of  a  fine  white  powder  when  con- 
centrated solutions  of  a  calcium  salt  and  some  sulfate  are 
mixed  together. 

Plaster  of  Paris,  (CaS04)2  •  H20.  This  is  a  fine  white 
powder  obtained  by  carefully  heating  gypsum.  When  water 
is  added  to  the  powder  a  plastic  mass  is  formed  which 
quickly  hardens,  or  sets.  This  property  makes  it  valuable 
for  molding  casts,  for  stucco  work,  and  for  a  finishing  coat 
on  plastered  walls.  Broken  bones  are  often  held  in  place 
by  casts  of  plaster  of  Paris  until  they  grow  together. 

In  the  manufacture  of  plaster  of  Paris  the  temperature  must 
not  be  allowed  to  rise  much  above  125°;  otherwise  the  anhy- 
drous salt  is  formed,  and  this  combines  with  water  so  slowly 
as  to  render  it  worthless  for  the  purposes  for  which  plaster 
of  Paris  is  used. 


THE  CALCIUM  FAMILY;  FERTILIZERS       431 

Hard  water.  Waters  containing  compounds  of  calcium 
and  magnesium  in  solution  are  called  hard  waters.  The 
hardness  of  water  may  be  of  two  kinds:  (1)  temporary 
hardness  and  (2)  permanent  hardness. 

1.  Temporary  hardness.    We  have  seen  that  when  water 
charged  with  carbon  dioxide  comes  in  contact  with  lime- 
stone a  certain  amount  of  the  latter  dissolves,  owing  to 
the  formation  of  the  soluble  acid  carbonate  of  calcium. 
The   hardness  of   such  waters  is  said  to  be  temporary, 
since  it  may  be  removed  by  boiling.     The  heat  changes 
the  acid  carbonate  into  the  insoluble  normal  carbonate, 
which  then  precipitates,  rendering  the  water  soft: 

Ca(HC03)2  — >•  CaC03  +  H2O  +  CO2 

Such  waters  may  also  be  softened  by  the  addition  of  suf- 
ficient lime  or  calcium  hydroxide  to  convert  the  acid 
carbonate  of  calcium  into  the  normal  carbonate: 

Ca(HCO3)2  -I-  Ca(OH)2  — >•  2  CaCO3  +  2  H2O 

2.  Permanent  hardness.    The  hardness  of  water  may  also 
be  due  to  the  presence  of  the  sulfate  or  the  chloride  of 
either  calcium  or  magnesium.    Boiling  the  water  does  not 
affect  these  salts;    hence  such  waters  are  said  to  have 
permanent  hardness.    They  may  be  softened,  however,  by 
the  addition  of  sodium  carbonate,  which  precipitates  the 
calcium  and  magnesium  as  insoluble  carbonates: 

CaS04  +  Na2C08  — ^  CaCO3  +  Na2SO4 
This  process  is  sometimes  called  "  breaking  "  the  water. 

Commercial  methods  for  softening  water.  The  average  water 
of  a  city  supply  contains  not  only  the  acid  carbonates  of  cal- 
cium and  magnesium  but  also  the  sulfates  and  chlorides  of 
these  metals,  together  with  other  salts  in  smaller  quantities. 


432    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Such  waters  are  softened  on  a  commercial  scale  by  the  addition 
of  the  proper  quantities  of  calcium  hydroxide  and  sodium  car- 
bonate. The  calcium  hydroxide  precipitates  the  acid  carbon- 
ates, while  the  sodium  carbonate  precipitates  the  other  soluble 
salts  of  calcium  and  magnesium.  The  amounts  of  calcium 
hydroxide  and  sodium  carbonate  required  to  soften  any  given 
water  are  calculated  from  a  chemical  analysis  of  the  water.  It 
will  be  noticed  from  the  equations  that  the  water  softened  in 

this  way  contains  sodium  sul- 
fate  and  sodium  chloride,  but 
the  presence  of  these  salts  is 
not  objectionable. 

Calcium  carbide  (CaC2). 
This  compound  is  prepared 
on  a  large  scale  for  use  in 
the  manufacture  of  acetylene 
(p.  302)  and  in  making  fer- 
tilizers. It  is  made  by  heat- 
ing a  mixture  of  lime  and 
coke  in  an  electric  furnace : 


FIG.  173.  A  furnace  for  the  manu- 
facture of  calcium  carbide 


CaO  +  3C 


The  pure  carbide  is  a  color- 
less, transparent  solid.    The 
commercial  article  is  a  dull- 
gray    porous    substance  which  contains  many  impurities. 
It  is  placed  on  the  market  in  air-tight  containers. 

Commercial  preparation.  While  calcium  carbide  was  first 
prepared  in  1836,  it  was  not  until  1893  that  it  became  a  com- 
mercial product.  The  general  principles  involved  in  its  prepara- 
tion are  illustrated  in  Fig.  173,  which  represents  a  simple  type 
of  carbide  furnace.  The  base  of  the  furnace  is  provided  with 
a  large  block  of  carbon  (^4),  which  serves  as  one  of  the  electrodes. 
The  other  electrodes  (B,  B~),  several  in  number,  are  arranged 
horizontally  at  some  distance  above  this.  A  mixture  of  coal 


THE  CALCIUM  FAMILY;  FERTILIZERS       433 

and  lime  is  fed  into  the  furnace  through  the  trap  top  C,  and 
in  the  lower  part  of  the  furnace  this  mixture  becomes  intensely 
heated  and  forms  liquid  carbide.  This  is  drawn  off  through 
the  tap  hole  D. 

Calcium  cyanamide  (CaCN2).  When  nitrogen  is  passed 
over  hot  calcium  carbide  the  two  react,  forming  a  com- 
pound known  as  calcium  cyanamide  : 


The  commercial  product  contains  about  60  per  cent  of  the 
cyanamide  ;  the  remaining  40  per  cent  consists  chiefly  of 
carbon  and  lime.  This  product  is  known  as  lime-nitrogen. 
This  is  ground  and  mixed  with  water  (which  slakes  the 
lime)  and  in  this  form  is  sold  as  a  fertilizer  under  the 
name  cyanamide.  Its  value  as  a  fertilizer  lies  in  the  fact 
that  all  of  its  nitrogen  is  available  as  a  plant  food. 

Calcium  cyanamide  promises  also  to  become  of  impor- 
tance in  the  commercial  preparation  of  ammonia  and  sodium 
cyanide,  both  of  which  can  readily  be  obtained  from  it  in 
a  manner  indicated  by  the  following  equations: 

CaCN2  +  3  H2O  —  >•  CaCO3  +  2  NH3 
CaCN2  +  C  +  2  NaCl  -  >-  CaCla  +  2  NaCN 

The  nitrogen  used  in  the  manufacture   of  cyanamide  is 
obtained  from  liquid  air  as  explained  on  page  131.    By 
means  of  the  reactions  expressed  in  the  above  equations 
it  is  possible,  therefore,  to  convert  the  nitrogen  of  the  air 
into  important  compounds  of  nitrogen. 

Phosphates  of  calcium.    With  phosphoric  acid,  calcium 
forms  three  salts,  the  names  and  formulas  of  which  are  as 
follows  : 

Normal  calcium  phosphate      ......     Ca3(PO4)2 

Calcium  monohydrogen  phosphate  ....     CaHPO4 

Calcium  dihydrogeii  phosphate    .....     Ca(H2l>O4)2 


434    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  normal  phosphate,  usually  called  simply  calcium 
phosphate,  is  found  in  quantities  in  nature,  largely  in  the 
form  of  rock  phosphate  or  phosphorite  and  as  a  constituent 
of  apatite.  It  is  the  chief  mineral  constituent  of  bones, 
the  ash  of  which  contains  about  80  per  cent  of  this  com- 
pound. The  phosphates  are  of  great  importance  in  con- 
nection with  the  subject  of  fertilizers  (see  p.  353  and  end 
of  this  chapter). 

Other  compounds  of  calcium.  Calcium  chloride  (CaCla)  occurs 
in  sea  water  and  is  formed  in  large  quantities  as  a  by-product 
in  the  Solvay  process  for  making  sodium  carbonate.  The  an- 
hydrous salt  readily  absorbs  moisture  and  is  used  as  an  agent 
for  drying  gases.  A  solution  of  the  salt  is  used  as  a  brine 
in  the  manufacture  of  ice  (p.  111).  It  has  also  been  used  to 
lay  the  dust  on  roads,  and  mines  have  been  sprinkled  with  it  in 
the  hope  of  preventing  dust  explosions.  Calcium  fluoride  (CaF2) 
occurs  in  nature  in  the  form  of  fluorite.  It  is  mined  in  large 
quantities,  especially  in  Illinois,  and  is  used  in  the  prepara- 
tion of  hydrofluoric  acid,  in  the  manufacture  of  opaque  glass, 
and  in  various  metallurgical  operations.  Calcium  sulfide  (CaS) 
is  a  by-product  in  the  Leblanc  process  for  making  sodium  car- 
bonate. The  commercial  salt  is  sometimes  used  as  a  luminous 
paint,  since,  after  exposure  to  a  bright  light,  it  will  glow  in 
the  dark.  Calcium  add  sulfite,  Ca(HS08)2,  is  used  as  a  preserv- 
ative, and  in  large  quantities  in  the  manufacture  of  paper 
(p.  330).  A  number  of  calcium  silicates  are  known  and  derive 
their  chief  interest  from  the  fact  that  they  are  important 
constituents  of  cement  and  glass. 

STRONTIUM  AND  BARIUM 

General.  These  elements  themselves  are  much  rarer  than 
calcium,  are  difficult  to  prepare,  and  have  no  commercial 
uses.  Their  most  abundant  minerals  are  the  following: 

Celestite SrSO4  Barite BaSO4 

Strontianite   ....     SrCO,          Witherite .  .     BaCO, 


THE  CALCIUM  FAMILY;  FERTILIZERS       435 

The  compounds  of  strontium  and  of  barium  are  similar 
in  composition  and  properties  to  the  corresponding  com- 
pounds of  calcium.  The  following  are  of  importance: 

Oxides  of  barium.  Barium  oxide  (BaO)  can  be  obtained 
by  strongly  heating  the  nitrate : 

2  Ba(NO3)2  — >•  2  BaO  +  4  NO2  +  O. 

When  heated  to  a  low  red  heat  in  the  air  the  oxide  com- 
bines with  oxygen,  forming  the  peroxide,  BaO2,  which  is 
used  in  making  hydrogen  peroxide  (p.  81). 

Barium  chloride  (BaCl2).  Barium  chloride  is  a  .white 
solid,  and  from  its  solution  it  crystallizes  as  the  hydrate 
BaCl2  •  2  H2O.  It  is  used  in  the  laboratory  as  a  reagent 
to  detect  the  presence  of  sulfuric  acid  or  soluble  sulfates, 
since  it  reacts  with  these  to  form  the  insoluble  barium 
sulfate. 

Barium  sulfate  (barite)  (BaSOJ.  Barium  sulfate  occurs 
in  nature  as  a  heavy  white  mineral  known  as  barite.  It  is 
precipitated  as  a  crystalline  powder  when  a  barium  salt 
is  added  to  a  solution  of  a  sulfate  or  to  sulfuric  acid: 

BaCl2  +  H2SO4  — *  BaSO4  +  2  HC1 

It  is  used  in  large  quantities  in  the  manufacture  of  paint. 

Barium  nitrate,  Ba(N03)2.  This  compound  is  an  oxidiz- 
ing agent.  When  ignited  with  some  oxidizable  substance 
it  gives  a  brilliant  green  light  and  is  used  for  this  purpose 
in  the  manufacture  of  fireworks. 

Strontium  hydroxide,  Sr(OH)2.  The  method  of  prepa- 
ration of  strontium  hydroxide  is  analogous  to  that  of 
calcium  hydroxide.  It  crystallizes  from  hot  water  in  the 
form  of  the  hydrate  Sr(OH)2 .  8  H2O.  Strontium  hydrox- 
ide has  the  property  of  forming  with  sugar  an  insoluble 
compound  which  can  easily  be  separated  again  into  its 


436    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


constituents.  It  has  therefore  been  used  in  the  sugar  refin- 
eries to  extract  sugar  from  impure  mother  liquors  from 
which  the  sugar  will  not  crystallize. 

Strontium  nitrate,  Sr(N03)2.  This  compound  imparts  a 
crimson  color  to  flames  and  is  used  in  the  manufacture  of 
fireworks  to  produce  a  red  light. 

FERTILIZERS 

Plant  food  ;  fertilizers.  With  the  exception  of  carbon  dioxide 
(and  possibly  a  little  oxygen)  absorbed  from  the  air,  the  grow- 
ing plant  derives  its  nourish- 
ment from  the  soil.  In  order 
that  vegetation  may  thrive  it 
is  essential,  therefore,  that  the 
soil  should  contain  an  ade- 
quate supply  of  appropriate 
plant  food.  Moreover,  since 
this  supply  is  continually  be- 
ing drawn  upon  by  the  grow- 
ing plant,  it  is  necessary,  in 
order  that  the  soil  may  retain 
its  fertility,  that  the  ingre- 
dients so  withdrawn  shall  be 
returned  to  it.  It  is  for  this 
purpose  that  fertilizers  are 
used. 

Constituents  of  fertilizers. 
While  a  number  of  elements 
are  essential  to  the  growth 
of  the  plant,  experience  has 
shown  that  in  general  the 
fertility  of  the  soil  may  be 

maintained  by  adding  three  substances :  (1)  nitrogenous 
matter,  (2)  phosphates  of  calcium,  and  (3)  compounds  of 
potassium.  Compounds  of  sulfur  are  also  essential,  but  they 
are  usually  present  in  sufficient  quantities. 


FIG.  174.  Justus  Liebig  (1803-1873) 

A  great  German  chemist  and  teacher. 

A  pioneer,  especially  in  agricultural 

chemistry  and  soil  fertility 


THE  CALCIUM  FAMILY;  FERTILIZERS       437 

Sources  of  fertilizers.  The  commercial  sources  of  each  of  the 
constituents  of  fertilizers  are  as  follows : 

1.  Nitrogenous  matter.    This  is  obtained  from  a  number  of 
sources :  sodium  nitrate,  ammonium  sulfate,  and  cyanamide ; 
also  nitrogenous  organic  matter,  such  as  dried  blood,  the  waste 
from  slaughterhouses,  and,  especially,  animal  excrements. 

2.  Phosphates.    Ground  bones  are  especially  valuable,  since 
they  contain  some  nitrogen  in  addition  to  calcium  phosphate. 
This  source,  however,  is   entirely  inadequate,  and  the  great 
supply  comes  from  the  rock  phosphates,  which  contain  about 
70  per  cent  of  calcium  phosphate.   These  rock  phosphates  are 
quarried  in  large  quantities,  especially  in  Florida  (Fig.  142) 
and  Tennessee.    Since  calcium  phosphate  is  nearly  insoluble, 
the  rock  is  ground  and  then  treated  with  sulfuric  acid.    This 
converts    the   insoluble    calcium   phosphate   into  the   soluble 
calcium  dihydrogen  phosphate,  Ca(H2P04)2: 

Ca3(P04)2  +  2  H2S04 >•  2  CaS04  +  Ca(H2P04)2 

The  resulting  mixture  of  calcium  sulfate  and  calcium  acid 
sulfate  is  a  powder  known  as  superphosphate  of  lime  and  is  a 
valuable  fertilizer.  The  calcium  sulfate  present  in  the  mixture 
adds  to  the  value  of  the  fertilizer,  furnishing  sulfur  and  im- 
proving the  physical  qualities  of  the  soil.  Certain  products 
(slags)  formed  in  the  manufacture  of  steel  contain  phosphorus 
and  are  used  in  fertilizers. 

3.  Potassium  compounds.   These  are  obtained  chiefly  from  the 
Stassfurt  mines.    Kainite  (KC1  •  MgS04  •  3  H20)  is  the  most 
common  of  the  minerals  used  (p.  408).    Wood  ashes  are  excel- 
lent, but  the  supply  is  limited. 

Commercial  fertilizers  are,  as  a  rule,  mixtures  of  the  three 
fundamental  materials  mentioned  above.  The  composition  of 
the  fertilizer  is  varied  according  to  the  crop  to  be  grown  and 
also  according  to  the  nature  of  the  soil  upon  which  the  fertilizer 
is  to  be  used 

Liming  soils.  Sometimes  a  soil  becomes  sour,  or  acid,  owing 
to  the  formation  of  acids  which  are  often  derived  from  decom- 
posing vegetable  matter.  Certain  plants,  such  as  mosses  and 


438    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

huckleberries,  will  thrive  in  acid  soil,  but  grass,  clover,  and 
grain  crops  will  not.  In  such  cases  the  soil  must  be  sweetened 
by  spreading  calcium  hydroxide  (slaked  lime)  upon  it  to  neu- 
tralize the  acids  present,  the  process  being  called  liming  the 
soil.  An  acid  soil  may  be  detected  by  moistening  strips  of 
blue  litmus  and  covering  them  for  a  few  minutes  with  the 
moist  soil. 

EXERCISES 

1.  What  properties  have  the  alkaline  earth  metals  in  common 
with  the  alkali  metals  ?    In  what  respect  do  they  differ  ? 

2.  Write  the  equation  for  the  reaction  between  calcium  carbide 
and  water. 

3.  How  is  calcium  chloride  removed  from  hard  water? 

4.  Would  air-slaked  lime   do  for  making  mortar?    Would  it 
serve  for  liming  acid  soils? 

5.  Why  would  you  expect  calcium  carbide  to  contain  impurities? 

6.  How  do  you  explain  the  fact  that  calcium  carbonate  can  be 
decomposed  into  calcium  oxide  and  carbon  dioxide,  and  yet  calcium 
oxide  absorbs  carbon  dioxide  from  the  air  to  form  the  carbonate? 

7.  Could  barium  hydroxide  be  used  in  place  of  calcium  hydrox- 
ide in  testing  for  carbon  dioxide  ? 

8.  Calcite  and  gypsum  often  resemble  each  other  in  appearance. 
How  could  yon  easily  distinguish  between  the  two  ? 

9.  Suggest  a  method  for  preparing  nitric  acid  from  calcium 
cyanamide. 

10.  Calcium  acid  sulfite  is  prepared  from  calcium  hydroxide  in 
a  manner  perfectly  analogous  to  the  preparation  of  the  acid  carbon- 
ate.   Write  the  equation  for  the  reactions  involved. 

11.  How  could  you  prepare  calcium  chloride  from  calcium  sul- 
fate  ?  barium  chloride  from  barite  ? 

12.  How  could  you  prove  that  dried  mortar  contains  calcium 
carbonate  and  sand? 

13.  Mention  different  advantages  gained  by  a  city  from  softening 
its  water  supply. 


THE  CALCIUM  FAMILY;  FERTILIZERS       439 

14.  What  weight  of  plaster  of  Paris  can  be  made  by  heating 
1  ton  of  gypsum  ? 

15.  What  weight  of  limestone  is  necessary  to  prepare  10  tons  of 
lime? 

16.  What  weight  of  water  is  necessary  to  slake  1  ton  of  lime  ? 

17.  Calculate  the  weight  of  calcium  oxide  present  in  the  lime 
made  by   heating  1  ton  of  limestone  containing  90  per  cent  of 
caloium  carbonate.    What  weight  of  water  would  be  required  to' 
slake  the  resulting  calcium  oxide  ? 

18.  The  heat  evolved  in  the  slaking  of  100  kg.  of  lime  would  raise 
the  temperature  of  what  weight  of  water  from  room  temperature 
(say  18°)  to  boiling  ? 

19.  A  certain  city  uses  10,000,000  gal.  of  water  daily,  and  100  gal. 
of  the  water  contains  120  g.  of  calcium  acid  carbonate  and  30  g.  of 
calcium  sulfate.    What  weights  of  calcium  hydroxide  and  sodium 
carbonate  are  required  to  soften  the  daily  water  supply  ? 

20.  In  the  manufacture  of  fertilizer  what  weight  of  sulfuric  acid 
containing  50  per  cent  by  weight  of  hydrogen  sulfate  woiild  be 
necessary  for  the  treatment  of  1  ton   of  phosphate  rock,  on  the 
supposition  that  the  only  reaction  taking  place  is  expressed  by 
the  following  equation : 

Ca3(P04)2  +  2H2S04 >-2CaS04  +  Ca(II2PO4)2 

21.  Starting  with,  limestone  and  sulfur,  how  could  you  prepare 
the  calcium  acid  sulfite  used  in  the  manufacture  of  paper? 


CHAPTER  XXXV 
THE  MAGNESIUM  FAMILY 


NAME 

SYMBOL 

ATOMIC 

WEIGHT 

DENSITY 

MELTING 

POINT 

BOILING 

POINT 

OXIDE 

Magnesium  .     . 
Zinc  .... 

Mg 
Zn 

24.32 
65.37 

1.74 
7.10 

651° 
419.4° 

920° 
950° 

MgO 
ZnO 

Cadmium     .     . 

Cd 

112.40 

8.64 

320.9° 

778° 

CdO 

The  family.  In  the  magnesium  family  are  included  the 
four  elements  magnesium,  zinc,  cadmium,  and  mercury. 
Between  the  first  three  of  these  metals  there  is  a  close 
family  resemblance,  such  as  has  been  traced  between  the 
members  of  the  two  preceding  families.  In  some  respects 
mercury  is  more  closely  related  to  copper  and  will  be 
studied  in  connection  with  that  metal. 

Properties.  At  ordinary  temperatures  oxygen  has  but 
little  action  upon  the  members  of  this  family.  At  high 
temperatures,  however,  combination  takes  place  rapidly, 
with  the  formation  of  oxides.  Magnesium  rapidly  decom- 
poses boiling  water,  while  zinc  and  cadmium  have  but 
slight  action  upon  it.  All  three  dissolve  in  acids,  with 
the  liberation  of  hydrogen. 

Compounds.  The  members  of  the  family  are  bivalent 
in  their  compounds,  so  that  the  formulas  of  their  salts 
resemble  those  of  the  alkaline  earth  metals.  Like  the 
latter  metals,  their  normal  carbonates,  normal  phosphates, 
and  normal  silicates  are  insoluble  in  water.  Their  sul- 
fates,  however,  are  readily  soluble.  Unlike  both  the  alkali 
440 


THE  MAGNESIUM  FAMILY  441 

metals  and  the  alkaline  earth  metals,  the  hydroxides  of  the 
metals  of  the  magnesium  family  are  nearly  insoluble  in 
water  and  are  much  more  readily  decomposed  by  heat, 
forming  water  and  the  oxide  of  the  metal. 

MAGNESIUM 

Occurrence.  Magnesium  is  a  very  abundant  element  in 
nature,  ranking  a  little  below  calcium  in  this  respect. 
Like  calcium,  it  is  a  constituent  of  many  rocks  and 
also  occurs  in  the  form  of  soluble  salts.  Dolomite 
(CaCO3  •  MgCO3)  and  magnesite  (MgCO3)  occur  in 
large  quantities.  Asbestos,  talc,  and  serpentine  are  sili- 
cates of  magnesium.  The  element  is  also  a  constituent 
of  chlorophyll,  the  green  coloring  matter  of  plants. 

Preparation.  The  metal  magnesium,  like  most  metals 
whose  oxides  are  difficult  to  reduce  with  carbon,  is  made 
by  electrolysis,  the  anhydrous  chloride  or  the  mineral 
carnallite  (p.  408)  being  used  as  the  electrolyte.  The 
electrolyte  is  melted  in  an  iron  pot,  which  also  serves  as 
the  cathode  in  the  electrolysis,  while  a  rod  of  carbon 
dipping  into  the  melted  salt  serves  as  the  anode.  The 
apparatus  is  very  similar  to  those  employed  in  the  prep- 
aration of  sodium  and  calcium.. 

Properties.  Magnesium  is  a  light,  silver-white  metal 
just  heavy  enough  to  sink  in  water.  It  is  usually  sold 
in  the  form  of  thin  ribbon  or  of  wire  or  as  a  powder. 
Air  does  not  act  rapidly  upon  it,  but  a  thin  film  of  basic 
carbonate  forms  upon  its  surface,  dimming  its  bright 
luster.  It  combines  directly  with  most  of  the  nonmetals, 
even  with  nitrogen.  The  common  acids  dissolve  it,  with 
the  formation  of  the  corresponding  salts.  It  can  be 
ignited  readily,  and  in  burning  it  gives  a  brilliant  white 


442    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

light.  This  light  is  very  rich  in  the  rays  which  affect 
photographic  plates,  and  the  metal,  in  the  form  of  fine 
powder,  is  extensively  used  in  the  production  of  flash 
lights  for  flash-light  photography,  for  white  lights  in  pyro- 
technic displays,  and  for  rockets  to  light  battlefields  by 
night.  When  used  for  this  purpose  the  powder  is  mixed 
with  an  oxidizing  agent,  potassium  chlorate  being  the  one 
commonly  employed. 

Magnesium  oxide  (magnesia)  (MgO).  Magnesium  oxide, 
sometimes  called  magnesia  or  magnesia  usta,  resembles 
lime  in  many  respects.  It  is  much  more  easily  formed 
than  lime  and  can  be  made  in  the  same  way ;  namely,  by 
heating  the  carbonate.  It  is  a  white  powder,  very  soft 
and  bulky,  and  is  unchanged  by  heat  even  at -very  high 
temperatures.  For  this  reason  it  is  used  in  the  manu- 
facture of  crucibles,  for  lining  furnaces,  and  for  other 
purposes  where  a  refractory  basic  substance  is  needed. 

Magnesium  hydroxide,  Mg(OH)2.  The  hydroxide  of 
magnesium  is  but  slightly  soluble  in  water  and  can  be 
precipitated  by  adding  a  soluble  base  to  a  salt  of 
magnesium : 

MgCl2  +  Ca(OH)2 >•  Mg(OH)2  +  CaCl2 

It  dissolves  sufficiently  to  give  a  slightly  alkaline  re- 
action and  is  a  moderately  strong  base.  It  is  a  white 
amorphous  substance  and  is  converted  into  the  oxide 
when  heated. 

Magnesium  carbonate  (MgC03).  Magnesium  carbonate 
occurs  in  a  number  of  localities  as  magnesite,  which  is 
usually  amorphous,  but  sometimes  forms  pure  crystals 
resembling  calcite.  More  frequently  it  is  found  associ- 
ated with  calcium  carbonate.  The  mineral  dolomite  has 
the  composition  CaCO8  •  MgCO3.  Limestone  containing 


THE  MAGNESIUM  FAMILY 


443 


smaller  amounts  of  magnesium  carbonate  is  known  as 
dolomitic  limestone.  Dolomite  is  one  of  the  most  common 
rocks,  forming  whole  mountain  masses.  It  is  harder  and 
less  readily  attacked  by  acids  than  limestone.  It  is  valu- 
able as  a  building  stone  and  for  foundations  and  as  ballast 
for  roadbeds.  Like  calcium  carbonate,  magnesium  car- 
bonate is  insoluble  in  water,  but  readily  dissolves  in  water 
containing  carbon  dioxide,  forming  the  acid  carbonate: 

MgC03  +  H20  +  C02  — >  Mg(HC03)2 

When  a  solution  of  a  magnesium  salt  is  precipitated 
with  sodium  carbonate,  there  is  obtained  a  white  basic 
carbonate,  known  as 
magnesia  alba,  which  is 
used  as  a  cosmetic  and 
as  a  medicine. 

Boiler  scale.  When 
water  which  contains 
certain  salts  in  solution 
is  evaporated  in  steam 
boilers,  a  hard  insoluble 
material  called  scale  de- 
posits in  the  boiler.  The 
formation  of  this  scale 
may  be  due  to  several 
distinct  causes : 

1.  To  the  deposit  of  cal- 
cium sulfate.  This  salt, 


FIG.  175.    Cross  section  of  a  boiler  tube 
showing  the  deposit  of  scale 


while  sparingly  soluble  in  cold  water,  is  almost  completely 
insoluble  in  superheated  water.  Consequently  it  is  precipitated 
when  water  containing  it  is  heated  in  a  boiler. 

2.  To  decomposition  of  add  carbonates.  As  we  have  seen,  cal- 
cium acid  carbonate  and  magnesium  acid  carbonate  are  decom- 
posed on  heating,  forming  insoluble  normal  carbonates  : 


Ca(HC08)2 


H2O  +  C02 


444    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

3.  To  hydrolysis  of  magnesium  salts.  Magnesium  chloride  and, 
to  some  extent,  magnesium  sulfate  undergo  hydrolysis  when 
superheated  in  solution,  and  the  magnesium  hydroxide,  being 
sparingly  soluble,  is  precipitated : 

MgCla  +  2  H20 ^Mg(OH)a  +  2  HC1 

The  scale  adheres  tightly  to  the  boiler  tubes  in  compact  layers 
(Fig.  175)  and,  being  a  nonconductor  of  heat,  causes  much 
waste  of  fuel.  It  is  very  difficult  to  remove,  owing  to  its  hard- 
ness and  its  resistance  to  reagents.  Thick  scale  sometimes 
cracks,  and  the  water  coming  in  contact  with  the  overheated 
iron  causes  an  explosion. 

Magnesium  sulfate  (MgSOJ.  Like  the  chloride,  magne- 
sium sulfate  is  found  rather  abundantly  in  springs  and 
in  salt  deposits.  Deposits  of  the  almost  pure  solid  salt 
having  the  composition  MgSO4  •  7  H2O  have  been  found 
in  Wyoming  and  Washington.  It  is  often  called  Epsom 
salt  because  of  its  occurrence  in.  the  waters  of  the  Epsom 
springs  in  England. 

Magnesium  sulfate  is  used  to  a  small  extent  in  the 
preparation  of  sodium  and  potassium  sulfates,  for  weight- 
ing cotton  cloth  in  the  dye  industry,  in  tanning,  and  in 
the  manufacture  of  paints  and  laundry  soaps.  To  some 
extent  it  is  used  in  medicine. 

Magnesium  silicates.  Many  silicates  containing  magne- 
sium are  known,  and  some  of  them  are  important  sub- 
stances. Serpentine,  asbestos,  talc  (or  soapstone),  and 
meerschaum  are  examples  of  such  substances.  Asbestos 
is  soft  and  fibrous  and  a  nonconductor  of  heat  and  of 
electricity.  It  is  used  for  fireproof  material  in  a  great 
variety  of  forms,  such  as  cloth,  paper,  board,  and  rope. 
It  is  also  used  as  a  covering  for  pipes,  furnaces,  and 
boilers,  to  diminish  heat  radiation.  It  also  has  many 
uses  as  an  insulator  in  electrical  devices.  The  chief 

t 


THE  MAGNESIUM  FAMILY  445 

source  of  asbestos  is  the  province  of  Ontario,  Canada. 
Soapstone  .is  valuable  for  sinks  and  table  tops  and,  in 
finely  ground  form,  as  a  toilet  powder  and  foot  ease.  It 
is  sometimes  called  French  chalk.  Meerschaum  is  used 
for  pipe  bowls  and  similar  articles. 


Occurrence.  Zinc  never  occurs  free  in  nature.  It  is  not 
a  constituent  of  common  rocks  and  minerals,  and  its 
occurrence  is  rather  local  and  is  confined  to  definite  de- 
posits or  to  pockets.  It  occurs  chiefly  in  the  following 
ores :  sphalerite  (zinc  blende)  (ZnS) ;  zincite  (ZnO) ; 
smithsonite  (ZnCO3)  ;  franklinite  (ZnO  •  Fe2O3)  ;  willem- 
ite  (Zn2SiO4).  One  fourth  of  the  world's  output  of  zinc 
comes  from  the  United  States  —  Missouri,  Kansas,  and 
New  Jersey  being  the  largest  producers. 

Metallurgy.  The  ores  employed  in  the  preparation  of 
zinc  are  chiefly  the  sulfide,  the  oxide,  and  the  carbonate. 
The  sulfide  and  the  carbonate  are  first  roasted  in  the  air, 
by  which  process  they  are  changed  into  the  oxide : 

ZnCO3 >-ZnO+CO2 

2  ZnS  +  3  O2 >-  2  ZnO  +  2  SO2 

The  oxide  is  then  mixed  with  coal  dust,  and  the  mixture 
is  heated  in  earthenware  retorts.  The  oxide  is  reduced 
by  this  means  to  the  metallic  state,  and  the  zirrc,  being 
heated  above  its  boiling  point,  distills  and  is  collected 
in  suitable  receivers  and  drawn  off  into  molds.  In  this 
form  it  is  called  spelter.  Commercial  zinc  often  contains 
impurities,  especially  carbon,  arsenic,  or  iron. 

Properties.  Pure  zinc  is  a  rather  heavy  bluish-white 
metal  with  a  high  luster.  It  melts  at  about  420°,  and  if 


446    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

heated  much  above  this  temperature  in  the  air,  it  takes  fire 
and  burns  with  a  bluish  flame.  It  boils  at  about  950°. 

Many  of  the  properties  of  zinc  are  much  influenced 
by  the  temperature  and  previous  treatment  of  the  metal. 
When  cast  into  ingots  from  the  liquid  state  it  becomes  at 
ordinary  temperatures  hard,  brittle,  and  highly  crystalline. 
At  100°-150°  it  is  malleable  and  can  be  rolled  into  thin 
sheets ;  at  higher  temperatures  it  again  becomes  very 
brittle.  When  once  rolled  into  sheets  it  retains  its  soft- 
ness and  malleability  at  ordinary  temperatures.  When 
melted  and  poured  into  water  it  forms  thin,  brittle  flakes, 
and  in  this  condition  is  called  granulated  zinc  or  mossy  zinc. 

Zinc  is  tarnished  superficially  by  moist  air,  but  beyond 
this  is  not  affected  by  it.  When  the  metal  is  quite  pure, 
sulfuric  acid  and  hydrochloric  acid  act  upon  it  very  slowly; 
when,  however,  it  contains  small  amounts  of  other  metals, 
such  as  magnesium  or  copper,  or  when  it  is  merely  in 
contact  with  another  metal  brisk  action  takes  place  and 
hydrogen  is  evolved.  For  this  reason,  when  pure  zinc  is 
used  in  the  preparation  of  hydrogen  a  few  drops  of  copper 
sulfate  are  often  added  to  the  solution  to  assist  the  chemi- 
cal action.  Strong  alkalies  dissolve  zinc,  liberating  hydrogen. 

Uses  of  zinc.  The  chief  use  of  zinc  is  in  the  manufacture 
of  galvanized  iron.  This  is  sheet  iron  or  wire  covered  with 
a  thin  layer  of  zinc,  which  protects  the  iron  from  rusting 
(p.  512).  About  two  thirds  of  all  the  zinc  produced  is 
used  in  this  way.  Sheet  zinc  is  used  as  a  lining  for  sinks 
and  water-containers.  Large  quantities  of  the  metal  are 
used  in  making  brass  and  other  alloys  (p.  497),  in  the 
construction  of  electrical  batteries,  and  in  separating  silver 
from  lead  (p.  516).  In  the  laboratory  it  is  used  in  the 
preparation  of  hydrogen  and,  in  the  form  of  zinc  dust,  as 
a  reducing  agent. 


THE  MAGNESIUM  FAMILY 


447 


Manufacture  of  galvanized  iron.  Fig.  176  shows  the  method 
used  in  making  galvanized  iron.  The  plates  of  iron  pass  under 
the  rollers  at  A  and  on  into  the  pot  of  melted  zinc  B.  The 
zinc  adheres  to  the  iron,  and  the  resulting  plate  is  passed  under 
the  roller  C  to  remove  the  excess  of  zinc  and  to  render  the 
surface  smooth.  Sometimes  the  zinc  is  deposited  on  the  iron 
by  electrolytic  methods. 

Zinc  oxide  (zinc  white)  (ZnO).  Zinc  oxide  occurs  in  im- 
pure form  in  nature,  being  colored  red  by  compounds  of 


FIG.  176.   The  manufacture  of  galvanized  sheet  iron 

manganese  or  of  iron.  It  can  be  prepared  in  the  same 
way  as  magnesium  oxide,  namely,  by  heating  zinc  car- 
bonate or  hydroxide,  but  it  is  more  often  made  by  burning 
the  metal. 

Zinc  oxide  is  a  pure-white  powder  which  is  much  used 
as  a  white  pigment  in  paints,  under  the  name  of  zinc 
white.  It  has  an  advantage  over  white  lead  in  that  it  is 
not  changed  in  color  by  sulfur  compounds,  while  lead 


448    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

compounds  turn  black.  Many  thousand  tons  of  zinc  oxide 
are  used  in  paints  each  year.  It  is  also  used  as  a  filler 
in  the  manufacture  of  rubber  goods.  Large  quantities  are 
used  annually  in  the  manufacture  of  automobile  tires. 

Zinc  sulfate  (ZnSOJ.  This  salt  is  readily  crystallized 
from  concentrated  solutions  in  transparent  colorless  crystals 
which  have  the  formula  ZnSO4  •  7  H2O  and  are  called  white 
vitriol.  It  is  prepared  commercially  by  the  careful  roasting 
ofthesulfide:  ZnS  +  2O 


Zinc  chloride  (ZnCl2).  This  salt  is  very  soluble  in  water 
and  has  a  strongly  acid  reaction.  It  has  germicidal  prop- 
erties and  is  used  to  preserve  railroad  ties  and  other 
timbers  especially  subject  to  decay. 

Zinc  sulfide  (ZnS).  This  substance  occurs  as  the  mineral 
sphalerite,  and  is  one  of  the  most  valued  ores  of  zinc. 
Very  large  deposits  occur  in  southwestern  Missouri.  The 
natural  mineral  is  found  in  large  crystals  or  masses  re- 
sembling resin  in  color  and  luster.  It  is  insoluble  in  water 
and  when  prepared  by  precipitation  is  white.  Lithopone 
is  a  mixture  of  the  two  solids  barium  sulfate  and  zinc 
sulfide  and  is  made  by  bringing  together  barium  sulfide 
and  zinc  sulfate  in  solution  : 

BaS  +  ZnSO4  -  >-  BaSO4  +  ZnS 
It  is  a  valuable  white  pigment. 

Preservation  of  wood.  With  the  rapid  disappearance  of  the 
forests  the  preservation  of  wood  from  decay  (fungous  growths) 
becomes  a  very  important  problem.  When  the  wood  is  to  be 
exposed  merely  to  atmospheric  conditions  it  is  preserved  by 
paints  and  varnishes.  When  it  must  be  partly  buried  in  the 
ground  (railway  ties,  fence  posts)  it  is  treated  with  germicidal 
preservatives.  Those  most  frequently  used  are  zinc  chloride, 
copper  sulfate,  and  coal-tar  creosote.__ 


THE  MAGNESIUM  FAMILY  449 

The  wood  is  placed  in  closed  boilers  in  baths  of  the  appro- 
priate liquid,  and  the  air  is  exhausted  so  that  the  liquid  may 
be  more  readily  driven  into  the  pores  of  the  wood.  Frequently 
the  latter  process  is  assisted  by  the  application  of  considerable 
pressure  to  the  liquid  after  the  air  has  been  pumped  out. 

CADMIUM 

The  element.  This  element  occurs  in  small  quantities  in 
some  zinc  ores.  In  the  course  of  the  metallurgy  of  zinc 
the  cadmium  compounds  undergo  chemical  changes  quite 
similar  to  those  of  the  zinc  compounds,  and  the  cadmium 
distills  along  with  the  zinc.  Being  more  volatile,  it  comes 
over  with  the  first  of  the  zinc  and  is  prepared- from  the  first 
portions  of  the  distillate  by  special  methods  of  purification. 
The  element  very  closely  resembles  zinc  in  most  respects. 
Some  of  its  alloys  are  characterized  by  having  low  melting 
points.  The  United  States  and  Germany  produce,  in  about 
equal  quantities,  the  world's  supply  of  cadmium. 

Compounds  of  cadmium.  Among  the  compounds  of  cad- 
mium may  be  mentioned  the  chloride,  CdCl2  •  2  H2Q,  the 
sulfate,  3  CdSO4  •  8  H2O,  and  the  nitrate,  Cd(NO3)2  •  4  H2O. 
These  are  white  solids  soluble  in  water.  The  sulfide,  CdS, 
is  a  bright-yellow  substance  which  is  insoluble  in  water  and 
in  dilute  acids.  It  is  valuable  as  a  pigment  in  fine  paints. 

EXERCISES 

1.  What  properties  have  the  metals  of  the  magnesium  family  iu 
common  with  the  alkali  metals  ?  with  the  alkaline  earth  metals  ? 

2.  Compare  the  action  of  the  metals  of  the  magnesium  group  on 
water  with  that  of  the  other  metals  studied. 

3.  What  metals  already  studied  are  prepared  by  electrolysis? 

4.  Write  the  equation  representing  the  reaction  between  mag- 
nesium   and    hydrochloric    acid ;   between    magnesium    and    dilute 
sulfuric  acid. 


450    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

5.  What  is  the  composition  of  commercial  lime  prepared  from 
dolomite  ? 

6.  With  phosphoric  acid  magnesium  forms  salts  similar  to  those 
of  calcium.    Write  the  names  and  formulas  of  the  magnesium  salts 
which  one  might  expect  to  be  thus  formed. 

7.  How  could  you  distinguish  between  magnesium  chloride  and 
magnesium  sulfate?  between  Glauber's  salt  and  Epsom  salt? 

8.  Account  for  the  fact  that  paints  made  of  zinc  oxide  are  not 
colored  by  hydrogen  sulfide. 

9.  Write  equations  showing  how  the  following  compounds  of 
zinc  may  be  obtained  from  metallic  /inc  :  the  oxide,  chloride,  nitrate, 
carbonate,  sulfate,  sulfide,  hydroxide. 

10.  How  could  you  prepare  the  pigment  cadmium  sulfide,  starting 
with  cadmium? 

11.  Note  the  position  of  the  alkali  metals,  the  calcium  family, 
and  the  magnesium  family  in  the  electrochemical  series.  How  would 
you  expect  acids  to  act  upon  these  metals  ? 

12.  Suggest  a  reason  why  none  of  the  metals  so  far  studied  are 
found  free  in  nature. 

13.  What  reaction  should  you  expect  to  take  place  if  a  piece  of 
zinc  is  immersed  in  a  solution  of  copper  sulfate  ?     Explain. 

14:  What  weight  of  carnallite  is  necessary  in  the  preparation  of 
500  g.  of  magnesium  ? 

15.  What  weight  of  franklinite  is  necessary  for  the  preparation 
of  1  ton  of  zinc  white  ? 

16.  1  g.  of  a  zinc  ore  was  dissolved  in  acid  and  the  zinc  present 
precipitated  by  hydrogen  sulfide.    The  resulting  zinc  sulfide  weighed 
0.38  g.     Calculate  the  percentage  of  zinc  in  the  ore. 

17.  Which  would  yield  the  more  zinc,  1  ton  of  sphalerite  or  1  ton 
of  f ranklinito  ? 


CHAPTER  XXXVI 
THE  ALUMINIUM  GROUP 

The  group.  With  the  exception  of  aluminium  none  of 
the  elements  of  Group  III  of  the  periodic  table  are  well 
known  or  abundant.  Boron  has  already  been  considered, 
and  the  others  fall  naturally  into  two  families.  The  one 
includes  aluminium,  together  with  gallium,  indium,  and 
thallium ;  the  other,  scandium  and  yttrium,  together  with 
a  large  group  of  elements  whose  oxides  are  collectively 
called  the  rare  earths. 

All  of  the  elements  of  this  group  are  trivalent  in  their 
compounds,  though  some  of  the  rarer  elements,  particularly 
thallium,  have  lower  valences  as  well.  With  few  excep- 
tions their  salts  are  colorless,  save  when  they  are  derived 
from  a  colored  acid.  The  bases  which  these  elements 
form  are  nearly  all  weak,  and  many  of  their  salts  are 
hydrolyzed  in  solution.  A  discussion  of  the  rare  elements 
of  the  group  does  not  fall  within  the  scope  of  an  ele- 
mentary text.  Aluminium,  however,  is  a  very  important 
metal,  so  that  it  will  be  considered  somewhat  in  detail. 

ALUMINIUM 

Occurrence.  Next  to  oxygen  and  silicon,  aluminium  is 
the  most  abundant  of  all  the  elements.  The  free  element 
is  not  found  in  nature,  but  its  compounds  are  widely 
distributed.  The  feldspars,  which  are  the  most  abundant 
of  all  the  minerals  in  the  earth's  crust,  are  all  silicates 
of  aluminium  and  either  sodium,  potassium,  or  calcium. 
451 


452    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


aluminium    oxides   (Al  O 


Since  the  soil  has  been  formed  largely  by  the  disintegra- 
tion of  these  rocks,  it  is  rich  in  the  silicates  of  aluminium, 
chiefly  in  the  form  of  clay.  Some  of  the  other  forms  in 
which  aluminium  occurs  in  nature  are  the  following :  corun- 
dum (A12O3)  ;  emery  (A12O8  colored  black  with  oxide  of 
iron)  ;  cryolite  (NagAlF6)  ;  bauxite,  a  mixture  of  hydrated 
H2O  and  A12O3 .  3  H2O)  to- 
gether with  similar  com- 
pounds of  iron.  Bauxite  is 
the  .ore  from  which  alumin- 
ium is  prepared.  In  the 
United  States  the  entire 
commercial  supply  now 
comes  from  Arkansas. 

Preparation.  Aluminium 
was  first  prepared  by  Wohler, 
in  1827,  by  heating  anhy- 
drous aluminium  chloride 
with  potassium : 


A1CL+3K 


8KC1+A1 


FIG.  177.    Charles  Martin  Hall 
1863-1914 

The  American  chemist  who  developed 
the  electrolytic  method  for  the  produc- 
tion of  aluminium 


Although  the  metal  is  very 
abundant  in  nature  and  pos- 
sesses many  desirable  prop- 
erties, the  cost  of  separating 
it  from  its  ores  was  so  great 
that  it  remained  almost  a 
curiosity  until  comparatively  recent  years.  With  the  de- 
velopment of  cheap  ways  of  obtaining  electrical  energy  the 
problem  has  been  solved,  and  the  metal  is  now  produced 
in  large  quantities  by  the  electrolysis  of  aluminium  oxide 
( A12O3)  dissolved  in  melted  cryolite  —  a  method  developed 
by  the  American  chemist  Hall  (Fig.  177)  in  1886. 


THE  ALUMINIUM  GROUP 


458 


Metallurgy.  An  iron  box  A  (Fig.  178)  about  8  ft.  long  and 
6  ft.  wide  is  connected  with  a  powerful  electrical  generator  in 
such  a  way  as  to  serve  as  the  cathode  upon  which  the  alu- 
minium is  deposited.  Three  or  four  rows  of  carbon  rods B,  B, 
dip  into  the  box  and  serve  as  the  anodes.  The  box  is  partly 
filled  with  cryolite,  and  the  current  is  turned  on,  generating 
enough  heat  to  melt  the  cryolite.  Aluminium  oxide  is  then 
added,  and  acts  as  an  electrolyte,  being  decomposed  into  alu- 
minium and  oxygen.  The  temperature  is  maintained  above  the 


FIG.  178.   Diagram  illustrating  the  manufacture  of  aluminium 

melting  point  of  aluminium,  and  the  liquid  metal,  being  heavier 
than  cryolite,  collects  on  the  bottom  of  the  vessel,  from  which 
it  is  tapped  off  from  time  to  time  through  the  tap  hole  C. 

Properties.  Aluminium  is  a  tin-white  metal  which  melts 
at  658.7°  and  is  very  light,  its  density  being  about  one 
third  that  of  iron.  It  is  stiff  and  strong,  and  with  frequent 
heating  can  be  rolled  into  thin  foil.*  It  is  a  good  con- 
ductor of  heat  and  electricity,  though  not  so  good  as 
copper^/br  a  given  cross  section  of  wire. 

Aluminium  is  not  perceptibly  acted  on  by  boiling  water, 
and  moist  air  merely  dims  its  luster.  Further  action  is 


454    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

prevented  in  each  case  by  the  formation  of  an  extremely 
thin  film  of  oxide  upon  the  surface  of  the  metal.  It  is  a 
good  reducing  agent,  combining  with  oxygen  at  high 
temperatures  with  the  liberation  of  much  heat : 

4  Al  +  3  Oo  — >-  2  A1,O8  +  760,400  cal. 

Hydrochloric  acid  acts  upon  it,  forming  aluminium  chlo- 
ride ;  nitric  acid  and  dilute  sulfuric  acid  have  almost  no 
action  on  it ;  but  hot  concentrated  sulfuric  acid  acts  upon 
it,  forming  aluminium  sulfate.  Alkalies  readily  attack  it, 
liberating  hydrogen : 

2  Al  +  6  KOH  — >-  2  A1(OK)3  +  3  H2 

Salt  solutions,  such  as  sea  water,  corrode  the  metal  rapidly. 

Uses  of  aluminium.  These  properties  suggest  many  uses 
for  the  metal.  Its  lightness,  strength,  and  inactivity  toward 
air  and  water  make  it  well  adapted  for  many  construction 
and  manufacturing  purposes.  These  same  properties  have 
led  to  its  extensive  use  in  the  manufacture  of  cooking 
utensils.  Owing  to  its  small  resistance  to  electrical  cur- 
rents, it  is  replacing  copper  to  some  extent  in  electrical 
construction,  especially  for  trolley  and  power  wires.  In 
the  form  of  a  powder  suspended  in  a  suitable  liquid  it 
makes  a  silvery  paint  used  to  cover  iron  pipes  and  lantern 
curtains.  The  greatest  use  of  aluminium  is  in  the  steel 
industry  (p.  482).  Aluminium  bronze,  consisting  of  about 
90  per  cent  copper  and  10  per  cent  aluminium,  has  a 
pure-golden  color,  is  strong  and  malleable,  is  easily  cast, 
and  is  permanent  ih  the  air.  Magnalium  is  an  alloy  of 
aluminium  and  magnesium.  It  is  light  and  rigid  and  is 
used  for  balance  beams. 

Goldschmidt  reduction  process.  Aluminium  is  frequently 
employed  as  a  powerful  reducing  agent,  many  metallic 


THE  ALUMINIUM  GROUP 


455 


oxides  which  resist  reduction  by  carbon  being  readily 
reduced  by  it.  The  aluminium,  in  the  form  of  a  fine 
powder,  is  mixed  with  the  metallic  oxide,  together  with 
some  substance  such  as  fluorite  to  act  as  a  flux.  The 
mixture  is  ignited,  and  the  aluminium  unites  with  the 
oxygen  of  the  metallic  oxide,  liberating  the  metal. 

Thermite  welding  process.  The  property  possessed  by  alu- 
minium of  reducing  oxides  with  the  liberation  of  a  large 
amount  of  heat  is  turned  into  practical 
account  in  the  welding  of  metals.  The 
German  chemist  Goldschmidt  was  the 
first  to  use  aluminium  for  this  purpose. 
The  welding  of  metals  by  this  method 
may  be  illustrated  by  a  single  example, 
namely  the  welding  of  car  rails  —  a 
process  often  carried  out  in  connection 
with  electric  railways  to  secure  good 
electrical  connection.  The  ends  of 
the  rails  are  accurately  aligned  and 
thoroughly  cleaned.  A  sand  mold  A 
(Fig.  179)  is  then  clamped  about  the 
ends  of  the  rail,  leaving  sufficient 
space  so  that  the  metal  can  flow  in. 
The  ends  of  the  rails  are  heated  to 
redness  by  the  flame  from  a  gasoline 

compressed-air  torch  directed  into  the  opening  in  the  mold. 
Just  over  the  opening  is  placed  the  conical-shaped  crucible  B, 
which  contains  a  mixture  of  iron,  metallic  oxides,  and  alu- 
minium. When  the  ends  of  the  rails  have  been  heated  ta  red- 
ness by  the  torch,  the  mixture  in  the  crucible  is  ignited,  and 
after  a  few  seconds  the  crucible  is  opened  at  the  bottom, 
and  the  molten  metal  resulting  from  the  reaction  in  the  crucible 
is  allowed  to  flow  into  the  mold.  The  molten  metal  surrounds 
the  ends  of  the  rails  and,  as  it  cools,  welds  them  firmly 
together.  A  mixture  of  the  metallic  oxides  and  aluminium 
ready  for  use  in  welding  is  sold  under  the  name  of  thermite. 


FIG.  179.  Welding  a  rail 
with  thermite 


456    AN  ELEMENTARY  STUDY  OF  CHEMISTKY 

Aluminium  oxide  (A1203).  This  substance  occurs  in  sev- 
eral forms  in  nature.  The  relatively  pure  crystals  are  called 
corundum ;  emery  is  a  variety  colored  dark  gray  or  black, 
usually  by  iron  compounds.  In  transparent  crystals,  tinted 
different  colors  by  traces  of  impurities,  it  forms  such 
precious  stones  as  the  sapphire,  ruby,  topaz,  and  oriental 
amethyst.  All  these  varieties  are  very  hard,  falling  little 
short  of  the  diamond  in  this  respect.  The  cheaper  forms, 
corundum  and  emery,  are  used  for  cutting  and  grinding 
purposes.  Chemically  pure  aluminium  oxide  can  be  made 
by  igniting  the  hydroxide,  when  it  forms  a  white  powder : 

2  Al(OH),  — *•  A1203  +  3  H20 

The  artificially  prepared  oxide  is  largely  used  in  the  prep- 
aration of  aluminium.  Some  laboratory  utensils,  such  as 
crucibles  and  tubes,  are  made  of  aluminium  oxide,  which 
is  given  the  trade  name  alundum.  The  same  material  is 
used  for  cutting  and  polishing  metals. 

Artificial  gems.  A  number  of  gems  are  now  prepared  in  the 
laboratory  from  molten  aluminium  oxide.  The  white  sapphires 
so  extensively  advertised  are  simply  the  pure  oxide.  By  incor- 
porating with  the  melted  oxide  small  percentages  of  certain 
metallic  oxides,  different  tints  or  colors  are  obtained,  and  in 
this  way  are  prepared  such  gems  as  the  ruby,  the  oriental 
amethyst,  and  the  yellow  sapphires  and  blue  sapphires,  all  of 
which  are  practically  identical  in  composition  and  properties 
with  the  natural  stones. 

Aluminium  hydroxide,  A1(OH)3.  The  hydroxide  can  be 
prepared  by  adding  ammonium  hydroxide  to  any  soluble 
aluminium  salt,  forming  a  colloidal  precipitate  which  is 
insoluble  in  water  but  very  hard  to  filter.  When  heated  it 
is  decomposed,  forming  the  oxide  and  water.  It  dissolves 


THE  ALUMINIUM  GROUP  457 

in  most  acids  to  form  soluble  salts,  and  in  the  strong  bases 
to  form  aluminates,  as  indicated  in  the  equations 

Al(OH).  +  3  HC1  — >•  A1C1.  +  3  H2O 
A1(OH)8  +  3  NaOH  — >-  Al(ONa)8  +  3  H2O 

It  may  act,  therefore,  either  as  a  weak  base  or  as  a  weak 
acid,  its  action  depending  upon  the  character  of  the  sub- 
stances with  which  it  is  in  contact.  It  is  therefore  an 
amphoteric  hydroxide  (p.  361). 

When  heated  gently  the  hydroxide  loses  part  of  its  hydrogen 
and  oxygen  according  to  the  equation 

A1(OH)3 >- A10  •  OH  +  H20 

This  substance,  the  formula  of  which  is  frequently  written 
HA102,  is  a  more  pronounced  acid  than  is  the  hydroxide,  and 
its  salts  are  frequently  formed  when  aluminium  compounds 
are  fused  with  alkalies.  The  magnesium  salt  Mg(A102)2  is 
called  spinel,  and  many  other  of  its  salts,  called  aluminabes, 
are  found  in  nature. 

Use  of  aluminium  hydroxide  in  water  purification.  The  value 
of  aluminium  hydroxide  in  the  purification  of  water  (p.  71)  is 
due  largely  to  its  colloidal  character.  This  colloid  is  dispersed 
all  through  the  water,  and  as  it  slowly  coagulates  and  settles 
it  carries  with  it  any  suspended  matter  present,  including 
microorganisms  and  coloring  materials.  The  colloidal  hydrox- 
ide is  formed  (as  hydrogel)  by  dissolving  in  the  water  some 
cheap  salt  which  readily  hydrolyzes,  such  as  aluminium  sulfate : 

A12(S04)8  +  6  H20 K  2  Al(OH),  +  3  H2S04 

As  a  rule  there  is  sufficient  basic  material  present  in  the  water 
to  combine  with  the  sulfuric  acid  set  free,  so  that  no  acid  is 
left  in  the  water  as  a  result  of  this  treatment ;  otherwise  some 
basic  substance  (such  as  calcium  hydroxide)  must  be  added 
with  the  aluminium  sulfate. 


458    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Fig.  180  illustrates  the  use  of  aluminium  sulfate  in  puri- 
fying water.  The  cylinder  A  contains  impure  water.  B  is 
a  similar  cylinder  of  water  to  which  some  aluminium  sul- 
fate has  been  added.  The  aluminium  hydroxide  formed 
by  hydrolysis  is  slowly  settling  in  the  water,  carrying  with 
it  the  impurities.  The  appearance  of  the  water  after  the 

aluminium  hydroxide   has 
settled  is  shown  in  C. 

Use  of  aluminium  hydroxide 
in  dyeing.  Aluminium  hy- 
droxide either  combines  with 
or  absorbs  many  soluble 
coloring  substances,  form- 
ing insoluble  products.  This 
property  leads  to  its  wide 
use  in  the  dyeing  industry. 
Most  of  the  dyes  are  pre- 
pared from  compounds  ob- 
tained from  coal  tar  (hence 
the  term  coal-tar  dyes}. 
Many  of  these  will  not  ad- 
here to  natural  fibers,  such 
as  cotton ;  that  is,  they  will  not  dye  fast.  It  is  often  pos- 
sible to  dye  such  cloth  in  the  following  way :  The  cloth  is 
first  soaked  in  a  solution  of  an  aluminium  salt.  It  is  then 
exposed  to  the  action  of  steam,  whereby  the  aluminium  salt  is 
completely  hydrolyzed,  the  resulting  aluminium  hydroxide  be- 
ing thus  thoroughly  incorporated  in  the  fiber.  If  the  cloth  is 
now  dipped  into  a  solution  of  the  dye,  the  aluminium  hydrox- 
ide combines  with  or  absorbs  the  color  substance  and  fastens, 
or  "  fixes,"  it  upon  the  fiber.  A  substance  such  as  aluminium 
hydroxide  which  serves  this  purpose  is  known  as  a  mordant. 

The  compounds  which  serve  well  as  mordants  may  be  pre- 
cipated  in  solutions  containing  various  dyes,  and  the  precipitate 
will  be  highly  colored,  though  not  always  of  the  same  color  as 
the  dye.  Colored  precipitates  of  this  kind  are  called  lakes. 


FIG.  180.    Purification  of  water  by 
aluminium  sulfate 


THE  ALUMINIUM  GROUP  459 

Aluminium  chloride  (A1C13).  The  anhydrous  chloride, 
which  is  used  in  the  synthesis  of  many  carbon  compounds, 
is  made  by  heating  aluminium  turnings  in  a  current  of 
chlorine.  The  hydrated  chloride,  A1C18  •  6  H2O,  is  pre- 
pared by  dissolving  the  hydroxide  in  hydrochloric  acid 
and  evaporating  to  crystallization.  When  heated  it  is 
converted  into  the  oxide,  resembling  magnesium  chloride 
in  this  respect: 

2  (A1C18  •  6  H20)  — >-  A1203  +  6  HC1  +  9  H2O 

Aluminium  sulfate,  A12(S04)3.  This  compound  is  pre- 
pared by  the  action  of  sulfuric  acid  011  bauxite  or  one  of 
the  common  silicates  of  aluminium.  It  crystallizes  from 
water  under  ordinary  conditions  as  A12(SO4)8  •  18  H2O.  It 
is  cheapest  of  the  soluble  salts  of  aluminium  and  is  used 
in  the  purification  of  water,  in  the  preparation  of  alums, 
and  in  certain  processes  connected  with  dyeing  or  with 
the  manufacture  of  paper. 

Alums.  Aluminium  sulfate  has  the  property  of  com- 
bining with  the  sulfates  of  the  alkali  .metals  to  form 
compounds  called  alums.  Thus,  with  potassium  sulfate  the 
reaction  is  expressed  by  the  equation 

K2S04  +  A12(S04)3  4-  24  H2O  — >•  2  (KA1  (SO4)2  - 12  H2O) 

The  sulfates  of  other  trivalent  metals  can  form  similar  com- 
pounds with  the  alkali  sulfates,  and  these  compounds  are  also 
called  alums,  though  they  contain  no  aluminium.  Tjhey  all 
crystallize  in  octahedra  and  contain  12  molecules  of  water  of 
hydration.  The  alums  most  frequently  prepared  are  the 
following : 

Potassium  alum KA1(SO4)2  •  12  H2O 

Ammonium  alum NH4A1(SO4)2  •  12  H2O 

Ammonium  iron  alum      ....     NH4Fe(SO4)2  •  12  H2O 
Potassium  chrome  alum  ....     IvCr(SO4)2  •  12  H2O 


460    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Very  large  well-formed  crystals  of  an  alum  can  be  prepared 
by  suspending  a  small  crystal  by  a  thread  in  a  saturated  solu- 
tion of   the  alum,  as   shown   in  Fig.  181. 
The  small  crystal  slowly  grows  and  often 
assumes  a  very  perfect  form. 

Hydrolysis  of  salts  of  aluminium.  While 
aluminium  hydroxide  forms  fairly  stable 
salts  with  strong  acids,  it  is  such  a  weak 
base  that  its  salts  with  weak  acids  are 
readily  hydrolyzed  (p.  226).  Thus,  when 
an  aluminium  salt  and  a  soluble  carbon- 
ate are  brought  together  in  solution,  we 
should  expect  to  have  aluminium  carbon- 
ate precipitated  according  to  the  equation 

3  Na2CO3  +  2  A1C13 >-  A12(CO8)3  +  6  NaCl 

But  if  it  is  formed  at  all,  it  instantly  begins  to  hydrolyze, 
the  products  of  the  hydrolysis  being  aluminium  hydroxide 
and  carbonic  acid,  the  latter  then  forming  carbon  dioxide 
and  water: 

2A1(OH)3  +  3H2C03 
3  H..O  +  3  CO0 


1 

SB 

FIG.  181.  Growing 

a  perfect   crystal 

of  alum 

3H2C03- 


It  is  because  of  these  reactions  that  alum  is  used  as  a 
constituent  of  some  baking  powders. 

Baking  powders.  Mixtures  of  sodium  bicarbonate,  flour  (or 
starch),  and  some  substance  that  will  act  upon  the  bicarbonate 
to  liberate  carbon  dioxide  are  known  as  baking  powders.  They 
are  used  as  aerating  agents  in  preparing  such  foods  as  cakes 
and  biscuits,  the  carbon  dioxide  evolved  pushing  its  way 
through  the  dough  and  rendering  it  porous  and  light.  The  com- 
pounds commonly  employed  for  liberating  the  carbon  dioxide 
from  the  sodium  bicarbonate  are  either  alum,  cream  of  tartar 


THE  ALUMINIUM  GKOUP  461 

(potassium  bitartrate),  or  calcium  hydrogen  phosphate,  and 
baking  powders  are  known  as  alum  baking  powders,  cream  of 
tartar  baking  powders,  or  phosphate  baking  powders,  according 
to  whether  they  contain  the  one  or  other  of  these  constituents. 
The  reactions  take  place  only  in  the  presence  of  water ;  hence 
the  use  of  the  flour,  which,  by  absorbing  any  moisture  that 
may  be  present,  prevents  the  powder  from  losing  its  strength 
until  used.  In  place  of  alum  a  mixture  of  sodium  sulfate  and 
aluminium  sulfate  known  as  cream  of  tartar  substitute,  or 
simply  as  C.  T.  S.,  is  largely  used.  The  complete  reaction  that 
takes  place  when  water  is  added  to  an  alum  baking  powder  is 
somewhat  complex,  but  it  amounts  to  the  formation  of  alumin- 
ium carbonate,  which  hydrolyzes  as  fast  as  formed,  evolving 
carbon  dioxide,  as  explained  in  the  preceding  paragraph. 

Aluminium  nitride  (A1N).  This  compound  is  prepared  by 
the  direct  union  of  aluminium  and  nitrogen  at  a  high 
temperature.  It  derives  its  chief  interest  from  the  fact 
that,  when  it  is  treated  with  steam,  ammonia  is  formed. 

2  A1N  +  3  H2O  — >•  AlaO3  +  2  NHg 

The  nitrogen  used  in  preparing  the  nitride  is  obtained 
from  the  air.  It  is  possible,  therefore,  through  the  inter- 
mediate formation  of  aluminium  nitride  to  convert  the 
nitrogen  from  the  air  into  ammonia. 

The  utilization  of  atmospheric  nitrogen.  Repeated  attempts 
have  been  made  to  utilize  the  inexhaustible  supplies  of 
free  nitrogen  present  in  the  atmosphere,  by  converting  the 
nitrogen  into  useful  compounds,  especially  ammonia,  nitric 
acid,  and  nitrates.  A  number  of  different  methods  for 
accomplishing  this  end  have  been  discussed  in  their 
appropriate  connections ;  it  is  desirable  to  collect  these 
together  under  one  heading. 

1.  The  nitrogen  may  be  converted  into  calcium  cyan- 
amide.  This  may  be  used  directly  as  a  fertilizer  or  it 


462    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

may  be   converted   into   ammonia   or   sodium   cyanide   as 
explained  on  page  433. 

2.  The  nitrogen  may  be  converted  into  ammonia.    This 
is  accomplished  by  heating  nitrogen  and  hydrogen  under 
the  proper  conditions  (p.  203). 

3.  The  nitrogen  may  be  converted  into  aluminium  nitride 
and  ammonia  prepared  from  this  as  described  above. 

4.  The  nitrogen  may  be  converted  into  nitric  acid.    This 
is  effected  by  sparking  mixtures  of  nitrogen  and  oxygen 
and  conducting  the  resulting  oxides  into  water  as  described 
on  page  208. 

At  present  it  is  cheaper  to  prepare  ammonia  from  coal 
and  nitric  acid  from  Chile  saltpeter.  Nevertheless,  it  is 
certain  that  increasing  amounts  of  these  compounds  will 
be  prepared  from  atmospheric  nitrogen  as  time  goes  on. 

EXERCISES 

1.  What  metals  and  compounds  studied  are  prepared  by  electrolysis? 

2.  Write  the  equation  for  the  reaction  between  aluminium  and 
hydrochloric  acid ;  between  aluminium  and  sulfuric  acid. 

3.  What  hydroxides  have  we  studied  that  have  both  acid  and 
basic  properties?    What  are  such  hydroxides  called? 

4.  Write    equations    showing   the   methods   used  for  preparing 
aluminium  hydroxide  and  aluminium  sulfate. 

5.  Write  the  general  formula  of  an  alum,  representing  an  atom 
of  a  univalent  metal  by  X  and  an  atom  of  a  trivalent  metal  by  Y. 

6.  What  is  the  significance  of  the  term  colloid  as  applied  to 
aluminium  hydroxide  ? 

7.  Compare  the  properties  of  the   hydroxides  of   the  different 
groups  of  metals  so  far  studied. 

8.  In  what  respects  does  aluminium  oxide  differ  from  calcium 
oxide  in  properties  ? 

9.  Where  should  you  expect  the  factories  for  the  production  of 
aluminium  to  be  located? 


THE  ALUMINIUM  GROUP  463 

10.  Why  do  the  directions  for  using  aluminium  cooking  utensils 
state  that  such  utensils  must  not  be  washed  in  alkaline  solutions? 

11.  Aluminium  sulfide  is  completely  hydrolized  by  water.    Write 
the  equations  for  the  reactions  you  would  expect  to  take  place  when 
a  solution  of  ammonium  sulflde  is  added  to  a  solution  of  aluminium 
chloride. 

12.  What  is  the  aerating  agent  used  in  the  making  of  bread? 
Why  not  use  baking  powder? 

13.  A  mixture  of  sour  milk  and  sodium  bicarbonate  is  sometimes 
used  as  an  aerating  agent  in  place  of  baking  powder.    What  is  the 
chemistry  involved? 

14.  What  is  the  derivation  of  the  word  mordantl    (Consult  dic- 
tionary.) 

15.  Give  four  different  ways  in  which  nitrogen  from  the  atmos- 
phere can  be  converted  into  nitric  acid. 

16.  1  kg.  of  potassium  alum  contains  what  weight  of  water  of 
hydration  ? 

17.  A  certain  bauxite  ore  was  found  to  contain  90  per  cent  of 

A1A 

1  ton  of  the  ore  ? 


CHAPTER  XXXVII 

ALUMINIUM   SILICATES  AND  THEIR  COMMERCIAL 
APPLICATIONS 

Aluminium  silicates.  One  of  the  most  common  constit- 
uents of  rocks  is  feldspar  (KAlSigOg),  a  mixed  salt  of 
potassium  and  aluminium  with  the  condensed  silicic  acid, 
H4Si3Og.  Under  the  influence  of  moisture,  carbon  dioxide, 
and  changes  of  temperature  this  substance  is  constantly 
being  broken  down  into  soluble  potassium  compounds  and 
aluminium  silicate  (Al2Si2O?  •  2  H2O).  In  relatively  pure 
condition  the  latter  is  called  kaolin  and  is  a  soft,  plastic 
mineral;  in  the  impure  state,  mixed  with  sand  and  other 
substances,  it  forms  common  clay.  Mica  is  another  very 
abundant  mineral,  having  a  varying  composition,  but 
being  essentially  of  the  formula  KAlSiO4.  Serpentine, 
talc,  asbestos,  and  meerschaum  are  important  complex  sili- 
cates of  aluminium  and  magnesium ;  granite  is  a  mechan- 
ical mixture  of  silica,  feldspar,  and  mica  and  is  therefore 
rich  in  aluminium.  Fuller's  earth  is  a  peculiar  form  of 
aluminium  silicate,  which  is  used  as  a  filtering  material 
for  decolorizing  oils,  especially  cottonseed  oil. 

Clay  products.  The  crudest  forms  of  clay  products, 
such  as  porous  brick  .and  draintile,  have  little  chemistry 
involved  in  their  manufacture.  Natural  clay  is  molded 
into  the  required  form,  dried,  and  then  burned  in  a  kiln, 
but  not  to  a  temperature  at  which  the  materials  soften. 
In  this  process  the  nearly  colorless  iron  compounds  in 
the  clay  are  converted  into  colored  compounds  which 
464  • 


ALUMINIUM  SILICATES 


465 


give  the  usual  red  color  to  these  articles.  In  making 
vitrified  brick  the  temperature  is  raised  to  the  point  at 
which  fusion  begins,  so  that  the  brick  is  partially  changed 
to  a  kind  of  glass. 

White  pottery.  This  term  is  applied  to  a  variety  of 
articles  ranging  from  the  crudest  porcelain  to  the  finest 
chinaware.  While  the  processes  used  in  the  manufacture 
of  the  articles  differ  in  details,  fundamentally  they  are 
the  same  and  may 
be  described  under 
three  heads;  name- 
ly, (1)  the  prepara- 
tion of  the  body  of 
the  ware,  (2)  the 
process  of  glazing, 
and  (3)  the  deco- 
ration. 

The  -body  of  the 
ware.  The  materials 
used  consist  of  clay 

artificially  prepared  FIG.  182.  The  manufacture  of  pottery :  mold- 
from  kaolin,  plastic  ing  the  plastic  material  into  form 

clay,  and  pulverized 

feldspar.  This  mixture  is  plastic  and  is  worked  into  the  desired 
shape  by  molds  or  on  a  potter's  wheel  (Fig.  182).  The  ware  is 
then  dried  and  burned  in  a  kiln  (Fig.  183)  until  vitrified,  and 
in  this  form  is  known  as  bisque.  This  is  usually  porous  and 
must  therefore  be  'glazed  to  render  it  nonabsorbent  and  give  it 
a  smooth  surface. 

The  glaze.  The  glaze  is  a  fusible  glass  which  is  melted 
over  the  surface  of  the  body.  The  constituents  of  the  glaze 
are  silica,  feldspar,  and  various  metallic  oxides,  often  mixed 
with  a  little  boric  oxide.  These  materials  are  finely  ground  and 
mixed  with  water  to  a  paste.  Sometimes  they  are  first  fused 
into  a  glass,  which  is  then  powdered  and  made  into  the  paste. 


466    AN  ELEMENTARY  STUDY  OF  CHEMISTEY 


The  bisque  is  dipped  into  the  glaze  paste,  dried,  and  fired  until 
the  glaze  materials  melt  and  flow  evenly  over  the  surface. 

The  decoration.  If  the  article  is  to  be  decorated,  the  design 
may  be  painted  upon  the  body  before  the  article  is  glazed,  or 
it  may  be  painted  upon  the  glaze  and  the  article  fired  again, 
the  pigments  melting  into  the  glaze.  In  the  former  case  the 
pigments  used  are  as  a  rule  metallic  oxides  of  various  colors, 

while  in  the  latter 
case  they  are  often 
colored  glasses. 

Cement.  The  term 
cement  as  ordinarily 
used  at  present  is 
applied  to  those 
mortars  which  pos- 
sess the  property  of 
hardening  in  water 
as  well  as  in  air. 
These  cements  are 
silicate  bodies,  usu- 
ally very  highly 
basic  in  character, 
and  when  ground 
fine  and  mixed  with 
water  they  undergo 
complex  reactions  resulting  in  the  formation  of  a  hard,  rock- 
like  mass.  A  number  of  different  classes  of  cements  are 
known,  the  most  important  of  which  is  called  Portland  cement. 
Composition  of  Portland  cement.  The  essential  ingredients 
of  Portland  cement  and  their  percentages  are  as  follows : 

SiO0 19  to  26%      MgO 0  to  5% 

A12O8 4  to  11%      S03 0  to  2.5% 

Fe203 2  to  5%        Na20|  _     Q  to  3% 

CaO 58  to  67%      K2O 


FIG.  183.  The  manufacture  of  pottery:  stack- 
ing the  ware  in  the  kiln  for  firing 


ALUMINIUM  SILICATES  467 

Manufacture  of  Portland  cement.  The  materials  most  com- 
monly employed  are  limestone  or  marl  and  clay  or  shale.  In 
general,  however,  any  substance  may  be  used  which  furnishes 
the  ingredients  listed  in  the  above  table.  Among  the  sub- 
stances so  used  is  blast-furnace  slag,  which  is  an  impure 
calcium-aluminium  silicate. 

The  materials  to  be  used  are  coarsely  ground  and  then 
mixed  together  in  the  proper  proportions  and  finely  pulver- 
ized. The  resulting  mixture  is  run  into  a  furnace  and  burned 


FIG.  184.   A  bridge  built  of  reenforced  concrete 

to  a  temperature  just  short  of  fusion,  at  which  temperature  it 
vitrifies,  forming  a  grayish  mass  known  as  clinker.  Finally, 
the  clinker  is  ground  to  a  fine  powder.  Gypsum  is  often  added 
in  the  process ;  this  acts  as  a  negative  catalyzer,  retarding  the 
hardening,  or  setting,  of  the  cement. 

The  setting  of  cement.  The  reactions  which  take  place 
upon  the  addition  of  water  to  cement  and  which  result 
in  the  formation  of  a  hard,  rocklike  mass  are  not  thor- 
oughly understood.  The  constituents  of  the  cement  ap- 
parently undergo  hydrolysis  when  they  come  in  contact 
with  water.  The  resulting  compounds  unite  with  water, 
producing  hydrates.  These  hydrates  are  crystalline  in 
character  and  form  a  hard,  compact  mass. 


468    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Growing  importance  of  cement.  Cement  is  rapidly  coming  into 
use  for  a  great  variety  of  purposes.  It  is  often  used  in  place 
of  mortar  in  the  construction  of  brick  buildings.  Mixed  with 
crushed  stone  and  sand  it  forms  concrete,  which  is  used  in 
foundation  work  for  buildings  and  street  paving.  It  is  also 
used  in  making  artificial  stone,  terra-cotta  trimmings  for  build- 
ings, artificial-stone  walks  and  floors,  fence  posts,  and  the  like. 
It  is  being  used  more  and  more  for  making  articles  which  were 
formerly  made  of  wood  or  stone,  and  the  entire  walls  of  build- 
ings are  sometimes  made  of  cement  blocks  or  concrete.  Iron 
rods  or  wire  are  often  embedded  in  the  concrete  before  it  sets, 
to  give  it  greater  strength,  and  this  is  called  reenforced  concrete. 


EXERCISES 

1.  In  the  manufacture  of  pottery  why  is  the  glaze  made  more 
fusible  than  the  body  of  the  ware? 

2.  Suppose  that  the  glaze  and  the  body  expand  and  contract  at 
different  rates  with  changes  in  temperatures,  what  will  be  the  result? 

3.  What  is  the  meaning  of  the  word  vitrify  ? 

4.  What  is  a  catalyzer  ?   What  is  a  negative  catalyzer  ? 

5.  Why  can  cement  be  used  as  jnortar  in  colder  weather  than 
ordinary  mortar? 

6.  What  is  meant  by  a  condensed  acid  ?   Give  example. 

7.  What  weight  of  kaolin  will  result  from  the  weathering  of 
1  ton  of  feldspar? 


CHAPTER   XXXVIII 
THE  IRON  FAMILY 


NAME 

SYMBOL 

ATOMIC 

WEIGHT 

DENSITY 

MELTING 

POINT 

OXIDES 

Iron   
Cobalt    

Fe 
Co 

55.84 
58.97 

7.86 
8.6 

1530° 
1480° 

FeO,  Fe20s 
CoO,  Co2O3 

Nickel    

Ni 

58.68 

8.9 

1452° 

NiO,  Ni203 

The  family.  The  elements  iron,  cobalt,  and  nickel  bear 
a  relation  to  one  another  which  is  different  from  that 
existing  among  the  members  of  any  other  family  as  yet 
considered.  Their  atomic  weights  are  very  close  together, 
and  in  the  periodic  table  they  are  placed  in  one  family 
not  because  the  plan  of  arrangement  brings  them  together 
but  because  they  are  so  similar  (p.  256)  and  evidently 
constitute  a  natural  family. 

The  elements  occur  in  nature  chiefly  as  oxides  and 
sulfides,  though  they  have  been  found  in  very  small 
quantities  in  the  native  state,  usually  in  meteorites. 
Of  the  three  iron  is  by  far  the  most  abundant,  cobalt  and 
nickel  being  of  rather  rare  occurrence.  Iron  and  nickel 
find  commercial  uses  chiefly  in  metallic  form,  while  cobalt 
is  more  widely  used  in  the  form  of  compounds.  Their  sul- 
fides, carbonates,  and  phosphates  are  insoluble  in  water, 
the  other  common  salts  being  soluble.  Their  salts  are 
usually  highly  colored,  those  of  iron  being  yellow  or  light 
green  as  a  rule,  those  of  nickel  darker  green,  while  cobalt 
salts  are  usually  rose  colored. 


470    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

IRON 

Occurrence.  The  element  iron  has  long  been  known, 
since  its  ores  are  very  abundant  and  it  is  not  difficult  to 
prepare  the  metal  from  them  in  fairly  pure  condition.  It 
occurs  in  large  deposits  as  oxides,  sulndes,  and  carbonates, 
and  in  smaller  quantities  in  a  great  variety  of  minerals. 
Indeed,  very  few  rocks  or  soils  are  free  from  small  per- 
centages of  iron.  It  is  a  constituent  of  the  chlorophyll 
of  plants  and  the  haemoglobin  of  the  blood  of  animals, 
and  therefore  plays  an  important  part  in  life  processes. 
Many  meteorites  are  largely  iron,  usually  alloyed  with 
a  little  nickel. 

Pure  iron.  Pure  iron  may  be  prepared  by  the  electrol- 
ysis of  a  solution  of  iron  sulfate  between  iron  electrodes, 
though  it  is  difficult  to  free  it  entirely  from  hydrogen  in 
this  way.  Hydrogen  from  the  water  of  the  electrolyte  is 
liberated  at  the  cathode  along  with  the  iron  and  is  dis- 
solved, or  occluded,  in  the  iron  as  the  latter  deposits. 
Iron  is  prepared  in  practically  pure  condition  by  the 
open-hearth  method  (p.  478).  It  is  a  silvery  metal  which 
melts  at  1530°.  It  is  ductile  and  malleable  and  almost 
as  soft  as  aluminium.  It  is  especially  well  adapted  to  the 
manufacture  of  electromagnets,  since  it  acquires  and  loses 
magnetic  properties  more  readily  than  do  the  ordinary 
varieties  of  iron.  It  is  also  used  for  purposes  where  resist- 
ance to  corrosion  is  desired,  for  it  does  not  rust  rapidly. 

The  iron  of  commerce.  Iron  differs  from  most  of  the 
other  metals  used  in  the  industries  in  that  the  pure  metal 
is  seldom  obtained  and  is  of  limited  application,  while 
that  containing  small  percentages  of  other  elements  ex- 
hibits a  wide  variety  of  properties  which  make  it  of  the 
greatest  value  for  many  different  purposes. 


THE  IRON  FAMILY  471 

Carbon  is  always  present  in  amounts  which  vary  from 
a  mere  trace  to  about  7  per  cent.  According  to  the  con- 
dition of  treatment  the  carbon  may  be  in  the  form  of 
graphite  scattered  through  the  iron,  or  it  may  occur  as 
a  solid  solution  of  carbon  in  iron,  or  as  carbides  of  iron. 
One  of  these  carbides  has  the  formula  Fe3C  and  is  called 
cementite.  Manganese,  silicon,  and  traces  of  phosphorus, 
sulfur,  and  oxygen  are  also  present. 

The  properties  of  the  iron  are  so  much  modified  by  the 
percentages  of  these  elements,  by  their  form  of  combina- 
tion, and  by  the  treatment  of  the  metal  during  its  produc- 
tion, that  many  varieties  of  iron  are  recognized  in  commerce, 
the  chief  of  which  are  cast  iron,  wrought  iron,  and  steel. 

The  metallurgy  of  iron.  The  problem  to  be  solved  in 
the  production  of  commercial  iron  is  (1)  to  obtain  a 
metallic  alloy  of  the  requisite  chemical  composition  and 
physical  properties  and  (2)  to  produce  it  on  a  very  large 
scale.  The  development  of  the  huge  modern  furnaces  has 
demanded  a  wonderful  application  of  chemical  knowledge 
to  a  definite  purpose  and  a  no  less  wonderful  engineering 
skill  in  securing  the  present  great  scale  of  production.  To 
understand  the  processes  to  be  described  it  will  be  neces- 
sary to  remember  constantly  that  large  and  rapid  produc- 
tion is  fully  as  necessary  as  great  purity. 

Materials  used  in  metallurgy  of  iron.  Four  different 
classes  of  materials  are  used  in  the  metallurgy  of  iron : 

1.  Iron  ore.  The  ores  most  frequently  employed  are  the 
following  : 

Hematite Fe2O3      Siderite  .     .     .     FeCO8 

Magnetite Fe3O4      Limonite     .     .     2  Fe2O3  •  3  H2O 

As    mined   for   use    all    ores   contain   earthy   matter  and 
often  sulfides  and  phosphates  as  well. 


472    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


While  iron  ore  is  mined  in  a  number  of  different  locali- 
ties in  the  United  States,  the  great  center  of  production 
is  in  the  neighborhood  of  Lake  Superior,  the  ore  being 
chiefly  hematite.  Fig.  185  represents  one  of  the  large 
mines  in  Minnesota.  Large  amounts  are  also  mined  near 
Birmingham,  Alabama. 

2.  Carbon.  Carbon  hi  some  form  is  necessary  both  as  a 
fuel  and  as  a  reducing  agent.  In  former  times  wood  char- 
coal was  used  to 
supply  the  carbon, 
but  now  coke  is 
almost  universally 
used. 

3.  Hot  air.  To 
maintain  the  high 
temperature  that  is 
required  for  the 
reduction  of  iron, 
a  very  active  com- 
bustion of  fuel  is 

necessary.  This  is  secured  by  forcing  a  strong  blast  of 
hot  air  into  the  lower  part  of  the  furnace  during  the 
reduction  process. 

4.  Flux.  All  the  materials  which  enter  the  furnace  must 
leave  it  again,  either  in  the  form  of  gases  or  as  liquids. 
The  iron  is  drawn  off  as  the  liquid  metal  after  its  reduc- 
tion, the  oxygen  with  which  it  was  combined  escaping  as 
oxide  of  carbon.  To  secure  the  removal  of  the  earthy 
matter  charged  into  the  furnace  along  with  the  ore,  ma- 
terials are  added  to  the  charge  which  will  combine  with 
the  impurities  in  the  ore,  forming  a  liquid.  The  material 
added  for  this  purpose  is  called  the  flux,  and  the  liquid 
produced  from  the  flux  and  the  ore  is  called  slag. 


FIG.  185.   Mining  iron  ore  in  Minnesota 


THE  IRON  FAMILY 


473 


The  slag.  The  slag  is  a  variety  of  difficultly  fusible 
glass,  being  essentially  a  calcium-aluminium  silicate.  If 
the  ore  is  rich  in  silica,  as  is  usual,  limestone  is  used  as  a 
flux ;  if  the  ore  contains  limestone,  silica  or  feldspar  is 
used;  if  the  ore  is  very  pure,  both 
constituents  must  be  added  as  flux. 

The  formation  of  slag  converts 
the  oxides  of  calcium,  magnesium, 
aluminium,  and  silicon  contained  in 
the  ore  into  the  liquid  state,  and  not 
only  does  this  make  the  removal  of 
these  materials  easy,  but  the  liquid 
is  a  necessity  for  other  reasons.  It 
is  a  medium  in  which  the  little  drop- 
lets of  iron  run  together  into  one 
large  mass  ;  it  keeps  the  contents  of 
the  furnace  fused  and  so  prevents 
clogging ;  it  floats  over  the  collected 
iron  and  prevents  its  oxidization. 
In  every  operation  in  which  iron  is 
melted  a  slag  must  be  provided. 

Cast  iron.  Ordinarily  the  first  step 
in  the  manufacture  of  any  variety 
of  commercial  iron  is  the  production 
of  cast  iron.  The  ores  are  mixed 
with  a  suitable  flux,  and  are  reduced  by  heating  with  coke. 

Blast-furnace  process.  The  reduction  is  carried  out  in  a  large 
tower  called  a  blast  furnace  (Fig.  186).  This  is  usually  80  ft. 
high  and  20  ft.  in  internal  diameter  at  its  widest  part,  narrow- 
ing somewhat  toward  both  the  top  and  the  bottom.  The  walls 
are  built  of  steel  and  are  lined  with  fire  brick.  The  base  is 
provided  with  a  number  of  pipes  (A~)  called  tuyeres,  through 
which  hot  air  is  forced  into  the  furnace.  The  tuyeres  are 


FIG.  186.  Vertical  section 
of  a  blast  furnace 


474    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

supplied  from  a  large  pipe  (5)  which  girdles  the  furnace.  At 
the  base  of  the  furnace  is  an  opening,  through  which  the  liquid 
metal  can  be  drawn  off  from  time  to  time.  There  is  also  a 
second  opening  (C)  somewhat  above  the  first,  through  which 
the  excess  of  slag  overflows.  The  top  is  closed  by  a  movable 
trap  (Z>)  called  the  bell,  and  through  this  the  materials  to  be 
used  are  introduced.  The  gases  resulting  from  the  combustion 
of  the  fuel  and  the  reduction  of  the  ore,  together  with  the 
nitrogen  of  the  air  admitted  through  the  tuyeres,  escape  through 


FIG.  187.    Casting  pig  iron  from  a  blast  furnace 

pipes  E.  These  gases  are  very  hot  and  contain  a  sufficient  per- 
centage of  carbon  monoxide  to  render  them  combustible ;  they 
are  accordingly  utilized  for  heating  the  blast  of  air  admitted 
through  the  tuyeres  and  as  fuel  for  the  engines. 

Charges  consisting  of  coke,  ore,  and  flux  in  proper  propor- 
tion are  at  intervals  introduced  into  the  furnace  through  the 
bell.  The  coke  burns  fiercely  in  the  hot-air  blast,  forming 
carbon  dioxide,  which  is  at  once  reduced  to  carbon  monoxide 
as  it  passes  over  the  highly  heated  carbon. 

Reduction  of  the  ore  begins  at  the  top  of  the  furnace  through 
the  action  of  the  •  carbon  monoxide.  As  the  ore  slowly  descends 
the  reduction  is  completed,  and  the  resulting  iron  melts  and 
collects  as  a  liquid  in  the  bottom  of  the  furnace,  the  lighter 


THE  IRON  FAMILY 


475 


slag  floating  above  it.  After  a  considerable  quantity  of  iron 
has  collected,  the  slag  is  drawn  off  through  C,  and  the  iron  is 
run  into  ladles  and  taken  to  the  steel  furnaces  for  the  manu- 
facture of  steel,  or  it  is  run  into  sand  molds  and  cast  into 
ingots  called  pigs.  Fig.  187  shows  the  method  of  drawing  off 
the  iron.  A  small  hole  is  made  near  the  bottom  of  the  furnace, 
and  the  molten  iron  flows  out  and  down  the  central  channel  A 
and  into  the  sand  molds  along  the  sides  B,  where  it  solidifies. 


FIG.  188.    A  typical  plant  for  the  manufacture  of  cast  iron 

In  practice  a  number  of  furnaces  are  usually  operated  to- 
gether, as  illustrated  in  Fig.  188,  which  shows  an  exterior  view 
of  a  modern  plant  for  making  cast  iron. 

Properties  of  cast  iron.  The  iron  produced  in  the  blast 
furnace  is  called  cast  iron.  It  varies  considerably  in  com- 
position, but  always  contains  over  2  per  cent  of  carbon, 
variable  amounts  of  silicon,  and  at  least  traces  of  phos- 
phorus and  sulfur.  The  form  in  which  the  carbon  is 
present,  whether  free  or  combined,  also  greatly  modifies 


476    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


the  properties  of  the  iron.  In  general,  cast  iron  is  hard 
and  brittle,  and  melts  at  about  1150°.  It  cannot  be 
welded  or  forged,  but  is  easily  cast  in  sand  molds,  it  is 
rigid,  but  not  elastic,  and  its  tensile  strength  is  small.  It 
is  used  for  making  castings,  but  chiefly  as  a  starting  point 
in  the  manufacture  of  other  varieties  of  iron. 

Wrought  iron.  Wrought  iron  is  made  from  cast  iron  by 
burning  out  most  of  the  carbon,  silicon,  phosphorus,  and 
sulfur,  the  operation  being  conducted 
in  what  is  called  a  puddling  furnace. 
Wrought  iron  is  soft,  malleable, 
and  ductile.  Its  tensile  strength  is 
greater  than  that  of  cast  iron,  but  less 
than  that  of  most  steel.  Its  melting 
point  is  much  higher  than  that  of  cast 
•iron.  If  melted,  it  is  changed  into 
steel.  It  is  no  longer  produced  to 
the  same  felative  extent  as  in  former 
years,  since  soft  steel  can  be  made  at 
a  less  cost  and  is  well  adapted  to 
almost  all  the  uses  for  which  wrought 
iron  was  used.  Its  chief  use  is  in 
making  materials  subject  to  corrosion,  such  as  water  pipes. 
Steel.  Steel,  like  wrought  iron,  is  made  from  cast  iron 
by  burning  out  a  part  of  the  carbon,  silicon,  phosphorus, 
and  sulfur,  but  the  processes  used  are  quite  different  from 
that  employed  in  the  manufacture  of  wrought  iron.  Nearly 
all  the  steel  of  commerce  produced  in  the  United  States  is 
made  by  one  of  two  general  methods  known  as  the  acid 
Bessemer  process  and  the  basic  open-hearth  process. 

Acid  Bessemer  process.  In  the  acid  Bessemer  process 
the  furnaces  used  are  lined  with  silica,  which,  it  will  be 
recalled,  is  an  acid  anhydride.  These  furnaces  remove 


FIG.  189.  Vertical  sec- 
tion showing  details  of 
a  Bessemer  converter 


THE  IKON  FAMILY 


477 


from  the  cast  iron  the  carbon  and  silicon,  but  not  the  phos- 
phorus and  sulfur.  The  process  is  therefore  employed  when 
the  cast  iron  to  be  used  is  low  in  phosphorus  and  sulfur. 

Details  of  operation  of  Bessemer  process.  This  process,  in- 
vented about  1880,  is  carried  out  in  great  egg-shaped  crucibles 
called  converters  (Fig.  189),  each  one  of  which  will  hold  as 
much  as  15  tons  of  steel.  The 
converter  is  built  of  steel  and 
lined  with  silica.  It  is  mounted 
on  trunnions,  so  that  it  can 
be  tipped  over  on  its  side  for 
filling  and  emptying.  One  of 
the  trunnions  is  hollow,  and  a 
pipe  connects  it  with  an  air 
chamber  (.4),  which  forms  a 
false  bottom  to  the  converter. 
The  true  bottom  is  perforated, 
so  that  air  can  be  forced  in  by 
an  air  blast  admitted  through 
the  trunnion  and  the  air 
chamber. 

White-hot  liquid  cast  iron 
from  a  blast  furnace  is  run 
into  the  converter  through  its 
open,  necklike  top  jB,  the  con- 
verter being  tipped  over  to  re- 
ceive it ;  the  air  blast  is  then 

turned  on,  and  the  converter  is  rotated  to  a  nearly  vertical  posi- 
tion. The  carbon  manganese  and  silicon  in  the  iron  are  rapidly 
oxidized  (first  the  silicon  and  manganese  and  then  the  carbon), 
the  oxidation  being  attended  by  a  brilliant  flame  (Fig.  190).  The 
heat  of  the  reaction,  largely  due  to  the  combustion  of  silicon, 
keeps  the  iron  in  a  molten  condition.  The  air  blast  is  continued 
until  the  character  of  the  flame  shows  that  all  the  carbon  has 
been  burned  away.  The  process  requires  on  the  average  about 
ten  minutes,  and  when  it  is  complete  the  desired  quantity  of 
carbon  (generally  in  the  form  of  high-carbon  iron  alloy)  is 


FIG.  190.   The  brilliant  flame  of 
Bessemer  converter 


478    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

added  and  allowed  to  mix  thoroughly  with  the  fluid.  The  con- 
verter is  then  tilted  and  the  steel  run  into  molds,  and  the 
ingots  so  formed  are  hammered  or  rolled  into  rails  or  other 
objects.  The  process  must  be  conducted  very  rapidly,  for  as 
soon  as  the  silicon  and  carbon  have  been  burned  there  is  no  way 
to  keep  the  iron  from  cooling,  and  it  must  be  poured  at  once. 

Basic  open-hearth  process.  In  the  basic  open-hearth 
process  the  lining  of  the  furnace  is  made  of  limestone  or 
dolomite,  both  of  which  act  as  bases.  In  such  furnaces 
the  phosphorus  and  sulfur  are  both  removed,  as  well  as 


"V 
HotAirJ     ^HotGas 

FIG.  191.    Vertical  section  of  an  open-hearth  furnace 

the  silicon  and  carbon.  The  presence  of  more  than  traces 
of  phosphorus  and  sulfur  in  the  finished  steel  renders  the 
metal  so  brittle  that  it  is  worthless  for  most  purposes. 
The  open-hearth  process,  therefore,  possesses  a  great  advan- 
tage over  the  acid  Bessemer  process  in  that  it  makes  it 
possible  to  utilize  iron  ores  (or  cast  iron  obtained  from 
them)  that  contain  appreciable  quantities  of  phosphorus 
and  sulfur.  The  operation  does  not  need  to  be  hastened, 
and  steel  of  any  desired  composition  can  be  produced. 

Details  of  the  open-hearth  process.  Fig.  191  shows  the  simpler 
parts  of  the  type  of  furnace  used  in  this  process.  The  hearth 
of  the  furnace  is  about  40  ft.  in  length,  12  ft.  in  width,  and 
2  ft.  in  depth,  and  is  lined  with  limestone  or  dolomite  (A,  A). 


THE  IRON  FAMILY  479 

Either  gas  or  sprayed  oil  is  used  as  fuel  and  finely  powdered 
coal  is  now  being  employed.  Below  the  furnace  is  placed  a 
checkerwork  of  brick  so  arranged  that  the  hot  products  of  com- 
bustion escaping  from  the  furnace  may  be  conducted  through 
it,  thus  heating  the  bricks  to  a  high  temperature.  Both  the  air 
necessary  for  combustion  and  the  gaseous  fuel  (unless  decom- 
posed by  heating,  as  in  the  case  of  natural  gas  and  sprayed  oil) 
are  preheated  by  passing  them  over  the  hot  bricks,  so  that  the 
temperature  reached  during  combustion  is  greatly  increased. 

The  gas  entering  through  C  comes  in  contact  at  D  with  the 
hot  air  entering  through  J3,  and  a  vigorous  combustion  results, 
the  flame  passing  above  and  over  the  cast  iron  and  lime  with 
which  the  furnace  is  charged.  The  products  of  combustion  es- 
cape through  E  and  F.  At  the  temperature  reached,  the  carbon 
in  the  cast  iron  is  removed  in  the  form  of  the  oxide,  the  escap- 
ing gas  giving  the  melted  metal  the  appearance  of  boiling.  The 
silicon,  phosphorus,  and  sulfur  unite  with  oxygen  to  form  acid 
anhydrides ;  these  combine  with  the  lime  to  form  a  slag,  and 
this  rises  to  the  surface  of  the  melted  charge  and  is  easily 
removed. 

When  a  test  shows  that  the  desired  percentage  of  carbon  is 
present  the  melted  steel  is  run  into  large  ladles  and  then  into 
molds.  An  average  furnace  produces  about  50  tons  of  steel  in 
a  given  charge,  approximately  eight  hours  being  required  in 
the  process.  At  present  by  far  the  largest  amount  of  steel 
produced  in  the  United  States  is  made  by  this  process. 

Tool  steel  or  crucible  steel.  Steel  designed  for  use  in  the- 
manufacture  of  edged  tools  and  similar  articles  should  be 
relatively  free  from  silicon  and  phosphorus,  but  should 
contain  from  0.5  to  1.5  per  cent  of  carbon.  The  percentage 
of  carbon  should  be  regulated  by  the  exact  use  to  which 
the  steel  is  to  be  put.  Steel-  of  this  character  is  usually 
made  in  small  lots  from  either  Bessemer  or  open-hearth 
steel  in  the  following  way. 

A  charge  of  melted  steel  is  placed  in  a  large  crucible 
and  the  calculated  quantity  of  pure  carbon  is  added.  The 


480    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

carbon  dissolves  in  the  steel,  and  when  the  solution  is 
complete  the  metal  is  poured  out  of  the  crucible.  This 
is  sometimes  called  crucible  steel. 

Electrothermal  metallurgy  of  steel.  An  increasing  quan- 
tity of  high-grade  tool  steel  is  being  produced  in  electrical 
furnaces.  The  furnace  is  heated  by  electrical  energy,  and 
a  large  quantity  of  steel  can  be  kept  melted  in  these  fur- 
naces as  long  as  may  be  desired.  The  electrical  current 
is  used  merely  to  produce  heat,  so  that  the  process  is  not 
dependent  upon  electrolysis.  This  method  is  almost  iden- 
tical with  the  open-hearth  method,  save  in  the  way  in 
which  the  heat  is  supplied,  and  produces  the  same  kind 
of  steel  as  does  the  latter  method. 

Properties  of  steel.  Steel  contains  from  a  trace  up  to 
2  per  cent  of  carbon,  less  than  1 0.1  per  cent  of  silicon, 
and  not  more  than  traces  of  phosphorus  and  sulfur.  When 
desired,  a  product  containing  as  high  as  99.85  per  cent 
of  iron  can  be  produced  by  the  open-hearth  method.  Such 
steel  is  very  soft,  but  resists  rust.  As  the  percentage  of 
carbon  increases,  the  steel  becomes  harder  and  less  ductile. 
Steel  can  be  rolled  into  sheets,  cast  in  molds,  and  forged 
into  desired  shapes. 

The  hardening  and  tempering  of  steel.  When  steel  con- 
taining from  0.5  to  1.5  per  cent  of  carbon  is  heated  to  a 
relatively  high  temperature  and  then  cooled  suddenly  by 
plunging  it  into  cold  water  or  oil,  it  becomes  very  hard 
and  brittle.  When  gradually  reheated  and  then  allowed 
to  cool  slowly  this  hardened  steel  becomes  softer  and  less 
brittle,  and  this  process  is  known  as  tempering. 

By  properly  regulating  the  temperature  to  which  the  steel  is 
reheated  in  tempering,  it  is  possible  to  obtain  any  condition  of 
hardness  demanded  for  a  given  purpose,  as  for  making  springs 
or  cutting-tools.  Steel  assumes  different  color  tints  at  different 


THE  IRON  FAMILY  481 

temperatures,  and  by  these  the  experienced  workman  can  tell 
when  the  desired  temperature  has  been  reached.  Lake  gives 
the  following  temperatures  for  the  tempering  of  tools : 

220° paper  cutters,  wood-engraving  tools 

240°' knife  blades,  rock  drills 

260° hand-plane  cutters  and  cooper's  tools 

275° axes,  springs 

290° needles,  screw  drivers 

300° wood  saws 

Alloy  steels.  As  we  have  seen  (p.  475),  small  quantities 
of  carbon  greatly  modify  the  properties  of  iron,  and  equally 
marked  effects  may  be  produced  by  a  great  many  other 
elements.  Accordingly,  to  secure  a  steel  with  the  requisite 
properties,  suitable  percentages  of  these  elements  are  added 
to  the  steel  just  before  it  is  run  out  of  the  furnace.  The 
elements  most  frequently  added  are  manganese,  silicon, 
nickel,  chromium,  tungsten,  vanadium,  and  titanium,  and 
steel  containing  an  appreciable  percentage  of  any  of  these 
elements  is  called  an  alloy  steel.  The  alloy  element  is 
added  in  the  form  of  a  rich  alloy  of  iron,  such  as  ferro- 
chromium  or  ferromanganese. 

The  approximate  percentage  of  metals  other  than  iron 
present  in  some  of  the  principal  steel  alloys,  as  well  as 
the  chief  uses  of  the  alloys,  is  as  follows : 

3.5%  nickel armor  plate,  nickel  steel 

3.5%  nickel  and  2.5%  chromium  .     .     .  armor  plate  and  projectiles 

12.0%  manganese  ...     .     ...     .     .  burglar-proof  safes 

5.0%  chromium  and  from  8  to  24% 

tungsten high-speed  lathe  tools 

0.1%  titanium car  rails  and  steel  castings 

0.2%  vanadium,  3.5%  nickel,  0.75% 

chromium  and  0.4%  manganese    .     .  automobile  springs  and  axles 

12-15%  silicon  (duriron  and  tantiron)  .  retorts  for  distilling  acids, 

electrodes 


482    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Steel  purifiers.  The  great  difficulty  in  securing  a  good 
steel  is  (1)  to  accomplish  the  complete  reduction  of  the 
oxide  and  (2)  to  prevent  the  absorption  of  gases  which 
cause  blowholes  as  the  casting  solidifies.  These  difficulties 
are  avoided,  as  far  as  possible,  by  adding  to  the  steel,  at 
the  close  of  the  operation,  certain  elements  which  will 
combine  with  the  oxygen  and  the  absorbed  gases.  The 
compounds  formed  pass  into  the  slag,  and  almost  none  of 
the  added  element  remains  in  the  finished  product.  Alu- 
minium is  used  to  a  large  extent  for  this  purpose,  as  are 
also  vanadium  and  titanium.  Such  elements  are  called 
purifiers  or  scavengers. 

Compounds  of  iron.  Iron  differs  from  the  metals  so  far 
studied,  in  that  it  is  able  to  form  two  series  of  compounds. 
In  the  one  series  the  iron  is  bivalent  and  forms  compounds 
which  in  formulas  and  many  chemical  properties  are  similar 
to  the  corresponding  zinc  compounds.  These  are  called 
ferrous  compounds.  In  the  other  series  iron  acts  as  a  tri- 
valent  metal  and  forms  salts  similar  to  those  of  aluminium. 
These  salts  are  known  as  ferric  compounds. 

Oxides  of  iron.  Iron  forms  several  oxides.  Ferrous  oxide 
(FeO)  is  not  found  in  nature,  but  can  be  prepared  arti- 
ficially in  the  form  of  a  black  powder  which  easily  takes 
up  oxygen,  forming  ferric  oxide : 

4  FeO  -|-  O2  — >•  2  Fe2O3 

Ferric  oxide  is  the  most  abundant  ore  of  iron  and  occurs 
in  great  deposits,  especially  in  the  Lake  Superior  region. 
It  is  found  in  many  mineral  varieties  which  vary  in  density 
and  color,  the  most  abundant  being  hematite,  which  ranges 
in  color  from  red  to  nearly  black.  When  prepared  artifi- 
cially it  is  a  bright-red  powder  which  is  used  as  a  pigment 
(Venetian  recT)  and  as  a  polishing  powder  (rouge). 


THE  IRON  FAMILY  483 

Magnetite  has  the  formula  Fe3O4  and  is  a  combination 
of  FeO  and  Fe2O3.  It  is  a  very  valuable  ore,  but  is  less 
abundant  than  hematite.  It  is  sometimes  called  magnetic 
oxide  of  iron,  or  loadstone. 

Ferrous  salts.  These  salts  are  obtained  by  dissolving 
iron  in  the  appropriate  acid  or,  when  they  are  insoluble, 
by  precipitation.  The  crystallized  salts  are  usually  light- 
green  in  color  and  are  hydrates.  Ferrous  hydroxide  is  a 
base  of  about  the  same  strength  as-  the  hydroxide  of  zinc 
or  of  magnesium.  Consequently  ferrous  salts  are  not  very 
much  hydrolized  in  solution. 

Ferrous  sulfate  (FeS04).  Ferrous  sulfate  is  the  most 
familiar  ferrous  compound.  It  is  usually  obtained  in  the 
form  of  the  hydrate  FeSO4  •  7  H2O,  called  copperas,  or  green 
vitriol,  and  is  prepared  commercially  as  a  by-product  in 
the  steel-plate  mills.  Preparatory  to  galvanizing  or  tin- 
ning (p.  447),  steel  plates  are  cleaned  by  immersing  them 
in  dilute  sulfuric  acid,  and  in  the  process  some  of  the  iron 
dissolves.  The  liquors  are  concentrated,  and  the  green 
vitriol  separates  from  them.  The  salt  is  used  in  the 
manufacture  of  ink  and  of  iron  alum,  as  a  substitute  for 
aluminium  sulfate  in  the  purification  of  water,  and  as  a 
reagent  to  destroy  weeds. 

Ferrous  sulfide  (FeS).  Ferrous  sulfide  is  sometimes  found 
in  nature  as  a  golden-yellow  crystalline  mineral  called 
pyrrhotite.  It  is  formed  as  a  black  precipitate  when  a 
soluble  sulfide  and  an  iron  salt  are  brought  together  in 
solution  : 


It  can  also  be  made  as  a  heavy  dark-brown  solid  by  fusing 
together  the  requisite  quantities  of  sulfur  and  iron.  It  is 
used  in  the  laboratory  in  the  preparation  of  hydrogen 
sulfide  (p.  233). 


484    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Iron  disulfide  (pyrite)  (FeS2).  This  substance  occurs  abun- 
dantly in  nature  in  the  form  of  brass-yellow  cubical  crystals 
and  in  compact  masses.  Sometimes  it  is  called  foots  gold 
from  its  superficial  resemblance  to  the  precious  metal.  It  is 
used  in  very  large  quantities  as  a  source  of  sulfur  dioxide 
in  the  manufacture  of  sulfuric  acid,  since  it  burns  readily 
in  the  air,  forming  ferric  oxide  and  sulfur  dioxide  : 

4FeS 


Ferrous  carbonate  (FeC03).  This  compound  occurs  in 
nature  as  siderite  and  is  a  valuable  ore.  Like  calcium 
carbonate,  it  dissolves  to  some  extent  in  water  containing 
carbon  dioxide,  forming  the  more  soluble  acid  carbonate  : 

FeCO3  +  H2C08  ^=±  Fe(HC08)2 

Spring  waters  containing  acid  carbonate  in  solution  are 
called  chalybeate  waters. 

Ferric  salts.  The  crystallized  ferric  salts  are  usually 
yellow  or  violet  in  color.  In  solution  and  in  the  absence 
of  free  acid  they  hydrolyze  even  more  readily  than  the 
salts  of  aluminium.  This  fact  indicates  that  ferric  hydroxide, 
Fe(OH)3,  is  a  very  weak  base. 

Ferric  hydroxide,  Fe(OH)3.  When  solutions  of  ferric  salts 
are  treated  with  ammonium  hydroxide,  ferric  hydroxide 
is  formed  as  a  rusty-red  precipitate  insoluble  in  water.  It 
might  be  thought  that  this  same  precipitate  would  form 
when  ferric  salts  undergo  hydrolysis  in  solution, 

FeCl8  +  3  H2O  —  >-  Fe(OH)8  +  3  HC1 

but  as  a  rule  no  precipitate  forms,  because  the  ferric  hy- 

droxide remains  in  the  form  of  a  colloid.    If  the  solution 

is  boiled,  some  of  this  colloid  is  coagulated  as  a  precipitate. 

Iron  rust  is  a  variable  mixture  of  hydrated  oxides  of 


THE  IRON  FAMILY  485 

iron.  When  a  film  of  rust  forms  on  iron  it  does  not  pro- 
tect the  metal  from  the  further  action  of  water  as  does  the 
rust  of  aluminium  and  zinc,  because  iron  rust  is  porous 
and  also  tends  to  scale  off,  exposing  a  fresh  surface. 

The  rusting  of  iron.  A  number  of  different  theories  have 
been  advanced  to  account  for  the  changes  taking  place  in  the 
rusting  of  iron.  The  most  satisfactory  of  these  is  known  as 
the  electrolytic  theory.  According  to  this  the  primary  reac- 
tion in  the  rusting  of  iron  is  between  iron  and  water,  as 
expressed  in  the  following  equation: 

Fe  +  2  (H+,  OH-)  -  >•  Fe++,  2  OH~  +  H2 

The  ions  Fe+  +  and  2  OH~  then  combine  to  form  ferrous  hy- 
droxide, Fe(OH)2.  This  is  further  acted  upon  by  oxygen  and 
water,  and  forms  the  complex  substance  known  as  iron  rust. 
It  is  evident  that  the  composition  of  rust  will  vary  according 
to  the  conditions  of  its  formation. 

Ferric  chloride  (FeCls).    This  salt  is  the  most  familiar 
of  the  ferric  salts.    It  can  be  obtained  most  conveniently 
by  dissolving  iron  in  hydrochloric  acid  and  then  passing 
chlorine  into  the  solution: 
Fe  +  2HCl 
2FeCl 


The  salt  when  crystallized  from  water  has  the   formula 
FeCls  •  6  H2O. 

Oxidation  of  ferrous  salts.  A  ferrous  salt  in  solution, 
exposed  to  the  action  of  the  air  or  of  an  oxidizing  reagent, 
is  rapidly  converted  into  a  ferric  salt  by  oxidation.  For 
example,  in  the  presence  of  hydrochloric  acid  the  oxidation 
of  ferrous  chloride  takes  place,  as  follows: 

2  FeCl2  +  [O]  +  2  HC1  —  *  2  FeCl8  +  H2O 

In  the  absence  of  free  acid  the  reaction  is  somewhat  more 
complicated,  but  is  of  the  same  order. 


486    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

It  will  be  seen  that  the  valence  of  the  iron  atoms  has 
been  increased  from  2  to  3.  This  increase  is  always  called 
oxidation  even  in  cases  in  which  no  oxygen  is  concerned, 
(p.  31)  as  is  illustrated  in  the  equation 

2  Fed.,  +  C12  —  >•  2  FeCl3 

This  latter  equation  may  be  written  in  such  a  way  as  to 
show  the  change  in  the  electrical  charges  upon  the  ions  : 

2(Fe++,  C1-,  Cl-)  +  Cla  —  >-2(Fe+  +  +,  C1-,  Cl~,  Cl~) 

The  charge  upon  each  cation  (Fe++)  has  been  increased 
from  2  to  3,  and  a  corresponding  number  of  anions  have 
been  formed.  In  general,  the  term  oxidation  is  applied  to  all 
reactions  in  which  the  valence  of  a  metallic  cation  is  increased. 
Reduction  of  ferric  salts.  The  changes  that  take  place 
when  a  ferric  salt  is  exposed  to  the  action  of  nascent 
hydrogen  or  other  reducing  agents,  such  as  metals,  are  the 
reverse  of  the  ones  just  described.  This  is  seen  in  the 
following  equations: 

FeCl3  +  [H]  —  >•  FeCl2  +  HC1 
2  FeCl8  +  Zn  -  >-  2  FeCl2  +  ZnCl2 

In  these  reactions  the  valence  of  the  iron  atoms  has  been 
lowered  from  3  to  2.  These  equations  may  also  be  written 
in  a  form  to  show  the  change  of  charge  upon  the  cations 


(Fe+  +  +,  C1-,  C1-, 

—  >•  (Fe++,  C1-,  C1-)  +  (H+,  C1-) 

Although  no  oxygen  has  been  removed,  the  ferric  chlo- 
ride is  said  to  be  reduced  (p.  48).  In  general  when  the 
valence  of  a  metallic  cation  is  diminished  the  salt  is  said 
to  be  reduced. 


THE  IRON  FAMILY  487 

Sodium  f errocyanide  (Na4FeC6N6) ;  potassium  f errocyanide 
(K4FeC6N6).  These  compounds  are  salts  of  the  unstable 
hydroferrocyanic  acid,  H4FeC6Ng.  They  are  prepared  from 
by-products  obtained  in  the  manufacture  of  coke.  When 
the  coal  is  heated  in  the  absence  of  air  small  amounts  of 
the  carbon,  nitrogen,  and  hydrogen  present  are  evolved 
in  the  form  of  hydrogen  cyanide.  This  is  absorbed  and 
converted  into  calcium  ferrocyanide  by  means  of  calcium 
hydroxide  and  the  spent  iron  oxide  employed  in  the 
purification  of  the  gas  evolved  in  the  coking  of  the  coal. 
The  calcium  ferrocyanide  so  obtained  is  converted  into 
sodium  or  potassium  ferrocyanide  by  treatment  with 
appropriate  salts  of  sodium  or  potassium.  The  reactions 
involved  in  the  complete  process  are  quite  complex. 

Both  sodium  ferrocyanide  and  potassium  ferrocyanide  are 
yellow  in  color  and  are  readily  soluble  in  water.  The  latter 
compound  is  often  called  yellow  pmssiate  of  potash.  The  so- 
dium salt  crystallizes  from  water  in  the  form  of  the  hydrate 
Na4FeC6Ng  •  10  H2O,  while  the  potassium  salt  crystallizes  as 
the  hydrate  of  K4FeC6N6  •  3  H2O.  In  solution  they  ionize  as 
follows :  Na^FeC^  — ^  4  Na+  +  FeC6N6- 
K4FeC6N6  — >-  4  K+  +  FeC6N6-  - 

It  is  important  to  notice  that  no  ions  of  iron  are  present, 
so  that  these  salts  do  not  give  the  ordinary  reactions  for 
iron.  They  react  with  ferric  salts  such  as  ferric  chloride 
in  accordance  with  the  following  equation : 

3  K4(FeC6N6)  +  4  FeCl8  — *  Fe4(FeC6N6)8  + 12  KC1 

The  resulting  ferric  ferrocyanide  is  a  blue  insoluble  solid 
and  is  ordinarily  called  Prussian  blue.  It  is  a  common 
paint  pigment,  and  most  of  the  ferrocyanides  made  are 
used  in  the  preparation  of  this  pigment. 


488    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Potassium  ferricyanide  (K3FeC6N6).  By  treating  a  solu- 
tion of  potassium  ferrocyanide  with  chlorine  water  and 
evaporating  the  solution  to  crystallization,  garnet-red  crys- 
tals are  deposited  which  have  the  composition  K3FeC6Na: 

2  K4FeC6N6  +  C12  — >•  2  K8FeC6N6  +  2  KC1 

This   compound  is   called  potassium  ferricyanide,   or  red 
prussiate  of  potash. 

In  place*  of  chlorine  one  may  use  nascent  oxygen 
evolved  by  electrolysis.  It  is  only  necessary  to  electrolyze 
a  solution  of  potassium  ferrocyanide  in  water.  The  oxygen 
evolved  at  the  anode  converts  the  potassium  ferrocyanide 
into  the  ferricyanide. 

Blue  printing.  When  a  ferric  salt  and  potassium  ferricyanide 
are  brought  together  in  solution  no  precipitate  forms,  though 
the  solution  acquires  a  yellowish  color.  On  exposure  to  the  sun- 
light the  ferric  salt  undergoes  a  partial  reduction  to  a  ferrous 
salt,  and  a  blue  precipitate  forms.  Advantage  is  taken  of  these 
facts  in  the  process  of  blue  printing.  A  sensitive  paper  is 
prepared  by  soaking  paper  in  a  solution  of  potassium  ferri- 
cyanide and  a  ferric  salt  (ferric  ammonium  citrate  is  generally 
used)  arid  drying  it  in  a  dark  place.  When  a  black  drawing 
on  tracing  cloth  is  placed  upon  such  a  sensitive  paper  and  the 
two  are  exposed  to  the  sunlight,  the  sensitive  paper  (except 
where  it  is  protected  by  the  black  lines)  turns  a  brownish  color. 
It  is  then  thoroughly  washed  with  water  to  remove  the  soluble 
salts,  during  which  process  the  portions  aeted  upon  by  the  light 
turn  blue,  while  the  unaffected  portions  are  left  white,  A  solu- 
tion of  sodium  hydroxide  can  be  used  as  an  ink  for  white  let- 
tering on  a  blue  print,  since  this  base  decolorizes  Prussian  blue. 

Other  salts  of  iron.  The  following  compounds  of  iron 
have  industrial  uses: 

Ferric  sulfate,  Fe2(SO4)3      ......     a  white  solid 

Ferric  nitrate,  Fe(NO3)8  •  6  H2O  ....     violet  crystals 

Iron  alum,  NH4Fe(SO4)2  •  12  H2O     .     .     .     violet  crystals 


THE  IRON  FAMILY  489 

Inks.  Many  of  the  common  black  inks  are  made  by  treating 
an  extract  of  nutgalls  with  ferrous  sulfate  and  adding  a  blue- 
black  dye.  The  nutgalls  are  rich  in  tannic  acid,  and  this,  with 
ferric  compounds  formed  by  the  oxidation  of  the  ferrous  sul- 
fate by  the  air,  gives  a  nearly  black  precipitate.  The  black  dye 
gives  a  temporary  color,  the  permanent  color  bjing  developed 
after  the  writing  has  been  exposed  to  the  air.  The  addition  of 
some  colloidal  material,  such  as  gum  arabic  or  dextrin,  together 
with  a  little  sulfuric  acid,  delays  the  precipitation  of  the  black 
substance  in  the  bottle.  A  preservative  is  usually  added  to 
prevent  the  ink  from  molding.  Formerly  all  black  inks  were 
made  by  the  above  method.  At  present  aniline  dyes  are  being 
used  more  and  more  for  making  such  inks. 

COBALT  AND  NICKEL 

Occurrence.  Cobalt  and  nickel  are  almost  always  found 
together  in  ores  which  also  contain  iron,  silver,  and  cop- 
per, in  combination  with  arsenic  and  sulfur.  The  richest 
deposits  are  in  Ontario  and  New  Caledonia.  The  extrac- 
tion of  these  metals  from  their  ores  and  their  separation 
from  each  other  is  too  complicated  a  process  to  be  de- 
scribed here.  Nickel  is  also  a  frequent  impurity  of  crude 
copper,  and  several  million  pounds  of  nickel  sulfate  are 
annually  recovered  in  the  United  States  in  the  refining 
of  copper  by  electrolysis. 

Properties  and  uses.  Both  these  metals  are  silvery  in 
appearance  and  take  a  high  polish.  They  are  somewhat 
heavier  than  iron,  and  melt  at  a  lower  temperature.  Their 
chief  use  is  in  making  alloys.  An  alloy  of  cobalt  and 
chromium  is  used  for  making  cutlery  and  lathe  tools. 
Nickel  coinage  consists  of  75  per  cent  of  copper  and  25  per 
cent  of  nickel.  German  silver  (p.  497)  also  contains  about 
25  per  cent  of  nickel.  Nickel  steel  usually  contains  about 
3.5  per  cent  of  nickel,  and  prior  to  the  war  more  than  half 


490    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

of  the  nickel  produced  was  used  in  making  this  product. 
Invar  contains  about  36  per  cent  of  nickel  and  does  not 
expand  when  heated.  Monel  metal  is  an  alloy  consisting 
of  60  per  cent  of  nickel,  33  per  cent  of  copper,  and  6.5  per 
cent  of  iron,  and  it  possesses  very  great  tensile  strength. 
Modern  rifle  bullets  are  made  of  an  alloy  of  copper  and 
nickel  of  about  the  composition  of  Monel  metal.  They 


FIG.  192.    The  process  of  electroplating  with  nickel 

are  as  hard  as  steel  and  much  heavier.  Nickel  is  exten- 
sively used  as  a  plating  upon  other  metals,  such  as  iron 
or  brass,  to  prevent  tarnishing  in  air,  and  cobalt  can  be 
used  in  the  same  way.  Finely  divided  nickel  is  used  as 
a  catalyzer  in  a  number  of  reactions,  such  as  the  conver- 
sion of  oils  into  solid  fats  (p.  343)  and  in  the  oxidation 
of  ammonia  into  nitric  acid. 

Electroplating  with  nickel.  Nickel  plating  is  accomplished 
by  an  electrolytic  process.  The  electrolyte  consists  of  a  solu- 
tion of  nickel  ammonium  sulfate,  a  salt  having  the  composition 


THE  IRON  FAMILY  491 

NiS04  •  (NH4)2S04  •  6  H20.  The  object  to  be  plated  is  sus- 
pended in  the  electrolyte  and  serves  as  the  cathode,  while  a 
plate  of  nickel  is  used  as  the  anode.  When  the  current  is  pass- 
ing through  the  electrolyte,  the  nickel  is  deposited  upon  the  ob- 
ject to  be  plated,  and  an  equivalent  portion  of  nickel  dissolves 
from  the  anode,  the  composition  of  the  electrolyte  remaining 
unchanged.  Fig.  192  illustrates  the  process  carried  out  on  a 
large  scale,  the  objects  to  be  plated  being  suspended  from  the 
rods  A,  A. 

Cobalt  oxide  (CoO).  This  is  the  form  in  which  most  of  the 
cobalt  comes  into  the  market.  It  is  a  black  powder  used  in 
making  other  cobalt  compounds  and  in  making  blue  glass  and 
blue  decorations  on  china.  Sometimes  ground  blue  cobalt  glass, 
called  smalt,  is  used  instead  of  the  oxide  and  as  a  pigment. 

Salts  of  cobalt  and  nickel.  Nearly  all  the  simple  salts  of 
cobalt  and  of  nickel  have  formulas  similar  to  those  of  ferrous 
salts.  The  most  familiar  are  the  following : 

Co(NO3)2  •  6  H2O a  cherry-red  deliquescent  salt 

CoCl2  •  6  H2O    .     .     ...-.'.     .     similar  in  appearance  to  the  nitrate 

CoS      .     .     .     .     .     ...     .     an  insoluble  black  precipitate 

NiSO4  •  7  H2O  ......     well-formed  green  crystals 

Ni(XO8)2  •  6  H2O deliquescent  green  crystals 

NiSO4  •  (NH4)2SO4  •  6  II2O     .     used  in  nickel  plating 

NiS an  insoluble  black  precipitate 


1.  In  the  manufacture  of  cast  iron,  why  is  the  air  heated  before 
being  forced  into  the  furnace  ? 

2.  Write  the  equations  showing  how  each  of  the  following  com- 
pounds of   iron  could  be  obtained  from  the  metal  itself :  ferrous 
chloride,  ferrous  hydroxide,  ferrous  sulfate,  ferrous  sulfide,  ferrous 
carbonate,  ferric  chloride,  ferric  sulfate,  ferric  hydroxide. 

3.  Account  for  the   fact  that  a  solution  of  sodium  carbonate 
when  added  to  a  solution  of  a  ferric  salt  precipitates  a  hydroxide 
and  not  a  carbonate. 

4.  Calculate  the  percentage  of    iron   in   each    of   the  common 
iron  ores. 


492    AN  ELEMENTAL Y  STUDY  OF  CHEMISTRY 

5.  Why  is  brass  often  nickel  plated? 

6.  Why  does  not  iron  occur  in  the  native  state  ?   What  does  its 
native  occurrence  in  meteorites  indicate  ? 

7.  Why  is  the  furnace  in  which  cast  iron  is  made  called  a  blast 
furnace  ? 

8.  Write  equations  for  the  oxidation  of  ferrous  sulfate  to  ferric 
sulf ate ;  for  the  reduction  of  ferric  sulfate  to  ferrous  sulfate. 

9.  Suggest  a  method  for  removing  the  iron  from  chalybeate 
waters. 

10.  Why  is  the  formula  for  Prussian  blue  written  Fe4(FeC6N6)3 
rather  than  Fe7(C6N6)3? 

11.  Will  ammonium  hydroxide  precipitate  ferric  hydroxide  when 
added  to  a  solution  of  sodium  f errocyanide  ? 

12.  AVhat  weight  of  iron  disulfide  is  necessary  for  the  preparation 
of  1000  kg.  of  50  per  cent  of  sulf  uric  acid  ? 

13.  One  ton  of  steel  prepared  by  the  Bessemer  process  is  found 
by  analysis  to  contain  0.2  per  cent  of  carbon.  What  is  the  minimum 
weight  of  carbon  which  must  be  added  in  order  that  the  steel  may 
be  made  to  take  a  temper? 


CHAPTER  XXXIX 
COPPER,  MERCURY,  AND  SILVER 


FORMULAS  OK 

KAJU 

SYMBOL 

ATOMIC 

DENSITY 

MELTING 

OXIDES 

-ous 

-lc 

Copper      .     .     . 

Cu 

63.57 

8.93 

1083.0° 

Cu20 

CuO 

Mercury    .     .     . 

Hg 

200.60 

13.56 

-  38.9° 

Hg20 

HgO 

Silver    .... 

Ag 

107.88 

10.50 

960.5° 

Ag20 

AgO 

The  family.  Although  mercury  is  not  in  the  same 
family  with  copper  and  silver,  the  three  elements  resemble 
each  other  so  closely  in  chemical  conduct  that  it  is  con- 
venient to  class  them  together  for  study.  Gold  belongs 
in  this  family,  but  will  be  described  later. 

Occurrence  and  properties.  The  three  elements  occur  in 
nature  to  some  extent  in  the  free  state,  but  are  usually 
found  as  sulfides.  Their  oxides  are  easy  to  reduce. 

All  three  are  heavy  metals  of  high  luster  and  are 
especially  good  conductors  of  heat  and  electricity.  They 
are  not  very  active  chemically.  They  occur  below  hydro- 
gen in  the  displacement  series,  and  as  a  consequence 
neither  hydrochloric  nor  dilute  su  If  uric  acid  has  any 
appreciable  action  upon  them.  Concentrated  sulfuric  acid 
attacks  all  three,  forming  metallic  sulfates  and  evolving 
sulfur  dioxide  (p.  247),  while  nitric  acid,  both  dilute  and 
concentrated,  converts  them  into  nitrates  with  the  evolu- 
tion of  oxides  of  nitrogen  (p.  210). 


494    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Two  series  of  salts.  Copper  and  mercury  form  oxides 
of  the  types  M2O  and  MO,  as  well  as  two  series  of  salts. 
In  one  series  the  metals  are  univalent  and  the  salts  have 
formulas  like  those  of  the  sodium  salts.  They  are  called 
cuprous  and  mercurous  salts.  In  the  other  series  the  metals 


FIG.  193.    Mining  copper  ore  at  Butte,  Montana 

are  bivalent  and  resemble  magnesium  salts  in  formulas. 
These  are  called  cupric  and  mercuric  salts.  Silver  forms 
only  one  series  of  salts,  being  always  a  univalent  metal. 


COPPER 

Occurrence.  The  element  copper  has  been  used  for 
various  purposes  since  the  earliest  days  of  history.  It  is 
often  found  in  the  native  state,  large  masses  of  it  occur- 
ring nearly  pure  in  the  Lake  Superior  region  and,  to  a 


COPPER,  MERCURY,  AND  SILVER     495 

smaller  extent,  in  other  places.    Apart  from  native  •  copper 
the  most  valuable  ores  are  the  following: 

SULFUR  ORES  OXYGEN  ORES 

Chalcopyrite      .     .     .  CuFeS2  Cuprite   .     .     Cu2O 

Chalcocite     ....  Cu2S  Melaconite  .     CuO 

Bornite Cu6FeS4  Malachite    .     CuCO3  •  Cu(OH)2 

Metallurgy  of  copper.  Ores  containing  little  or  no  sulfur 
are  easy  to  reduce.  They  are  first  crushed  and  the  earthy 
impurities  washed  away.  The  concentrated  ore  is  then 
mixed  with  carbon  and  a  flux  and  heated  in  a  furnace, 
metallic  copper  resulting 
from  the  reduction  of 
the  copper  oxide  by  the 
hot  carbon. 


Metallurgy  of  sulfur  ores. 
Much   of  the    copper   of 
commerce  is  made   from 
Chalcopyrite  and  bornite,      Flo-  194>  vertical  section  of  a  reverber- 
and  these  ores  are  more  atory  furnace 

difficult   to   work,  partly 

because  they  are  sulfides  and  cannot  be  reduced  by  heating 
with  carbon,  and  partly  because  they  contain  a  great  deal 
of  iron  as  well  as  copper. 

If  the  ore  is  in  fine  grains  it  is  first  smelted  in  a  reverbera- 
tory  furnace  (Fig.  194).  This  is  a  long  and  rather  narrow 
furnace  with  a  concave  floor  (A)  on  which  the  ore  B,  together 
with  silicious  flux,  is  placed  through  a  row  of  openings  along 
the  sides  of  the  roof.  The  ore  is  heated  by  a  long  flame  from 
burning  coal  led  through  the  furnace,  the  heat  from  the  flame 
being  reflected  down  upon  the  charge  from  the  arched  roof  C. 
Sometimes  the  flame  is  maintained  by  blowing  powdered  coal 
through  the  furnace. 

A  part  of  the  sulfur  is  burned  and  a  part  of  the  iron  is  con- 
verted into  iron  oxide,  which  then  combines  with  the  silica  to 


496    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

make  a  slag  of  ferrous  silicate.  The  copper  sulfide,  together 
with  the  remaining  iron  sulfide,  melts  to  form  a  heavy  liquid 
called  matte,  which  collects  under  the  slag. 

If  the  ore  is  coarse,  so  that  an  air  blast  will  not  blow  it 
away,  the  matte  is  made  in  a  blast  furnace  resembling  the  one 
used  for  making  cast  iron,  only  much  smaller. 

The  liquid  matte  is  then  tapped  off  into  a  converter  closely 
resembling  the  one  used  in  the  Bessemer  process  and  holding 
sometimes  as  much  as  50  tons.  Some  silica  is  added  and  a  care- 
fully regulated  current  of  air  is  blown  in.  The  sulfur  acts  as 
fuel,  burning  to  form  sulfur  dioxide ;  the  iron  is  converted 
into  oxide,  which  then  combines  with  the  silica  to  form  slag ; 
the  copper  sulfide  burns  to  form  sulfur  dioxide  and  copper. 
When  the  process  is  complete  the  liquid  copper  is  drawn  off 
and  poured  into  molds.  It  is  called  blister  copper  and  may  be 
as  high  as  98  per  cent  pure.  The  United  States  produces  an 
enormous  tonnage  of  copper  annually. 

Refining  of  copper.  Blister  copper  is  purified  by  electrol- 
ysis. A  large  plate  of  it,  serving  as  an  anode,  is  sus- 
pended in  a  tank,  facing  a  thin  plate  of  pure  copper  which 
is  the  cathode.  The  tank  is  filled  with  a  solution  of 
copper  sulfate  and  sulfuric  acid  to  act  as  the  electrolyte. 
A  current  from  a  dynamo  passes  from  the  anode  to  the 
cathode,  and  the  copper,  dissolving  from  the  anode,  is 
deposited  upon  the  cathode  in  pure  form,  while  the  im- 
purities collect  on  the  bottom  of  the  tank.  Electrolytic 
copper  is  one  of  the  purest  of  commercial  metals. 

Recovery  of  gold  and  silver.  Gold,  silver,  and  nickel  are  often 
present  in  small  quantities'  in  copper  ores,  and  remain  in  the 
erude  copper.  In  electrolytic  refining  the  gold  and  silver  col- 
lect in  the  muddy  deposit  on  the  bottom  of  the  tank.  The  mud 
is  carefully  worked  over  from  time  to  time  and  the  precious 
metals  extracted  from  it.  A  surprising  amount  of  gold  and 
silver  is  obtained  in  this  way.  The  nickel  passes  into  solution 
and  is  recovered  from  the  electrolyte  (p.  489). 


COPPER,  MERCURY,  AND  SILVER  497 

Properties  of  copper.  Copper  is  a  rather  heavy  metal,  of 
density  8.9,  and  has  a  characteristic  reddish  color.  It  is 
rather  soft  and  is  very  malleable,  ductile,  and  flexible,  yet 
tough  and  strong;  it  melts  at  1083°  and  boils  at  2310°. 
As  a  conductor  of  heat  and  electrical  energy  it  is  second 
only  to  silver. 

Since  it  is  below  hydrogen  in  the  displacement  series, 
hydrochloric  acid,  dilute  sulfuric  acid,  and  fused  alkalies 
are  almost  without  action  upon  it ;  nitric  acid  and  hot 
concentrated  sulfuric  acid,  however,  readily  dissolve  it 
(pp.  210,  247).  In  moist  air  it  slowly  becomes  covered 
with  a  film  of  the  bright-red  oxide  Cu2O,  which  soon 
changes  to  a  green  basic  carbonate.  If  heated  in  the  air 
the  metal  is  easily  oxidized  to  the  black  oxide  CuO. 

Uses.  Copper  is  extensively  used  for  electrical  purposes, 
for  roofs  and  cornices,  for  sheathing  the  bottoms  of  ships, 
and  for  making  alloys.  In  the  following  table  the  compo- 
sition of  some  of  these  alloys  is  indicated : 

Aluminium  bronze  .     .  90%~98%  copper,  2%~10%  aluminium 

Brass 63%-73%  copper,  27%-37%  zinc 

Bronze 70%-95%  copper,  l%-25%  zinc,  1%-18%  tin 

German  silver      .     .     .  50%~60%  copper,  20%  zinc,  20%-30%  nickel 

Gun  metal 90%  copper,  10%  tin 

Gold  coin 10%  copper,  90%  gold 

Silver  coin 10%  copper,  90%  silver 

Nickel  coin      ....  75%  copper,  25%  nickel 

Electrotyping.  Books  are  often  printed  from  electrotype 
plates,  which  are  prepared  as  follows :  The  face  of  the  type 
is  covered  with  wax,  and  this  is  firmly  pressed  down  until  a 
clear  impression  is  obtained.  The  impressed  side  of  the  wax 
is  coated  with  graphite,  and  this  is  made  the  cathode  in  an  elec- 
trolytic cell  containing  a  copper  salt  in  solution.  The  copper  is 
deposited  as  a  thin  sheet  upon  the  letters  in  wax  and,  when 
detached,  is  a  perfect  copy  of  the  type,  the  under  part  of  the 


498    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

letters  being  hollow.  The  sheet  is  strengthened  by  pouring  on 
the  undersurface  a  suitable  amount  of  commercial  lead.  The 
sheet  so  strengthened  is  then  used  in  printing. 

Two  series  of  copper  compounds.  Copper,  like  iron, 
forms  two  series  of  compounds :  the  cuprous  compounds,  in 
which  it  is  univalent ;  and  the  cupric  compounds,  in  which 
it  is  bivalent.  The  cupric  salts  are  much  the  more  common 
of  the  two. 

Cuprous  compounds.  The  most  important  cuprous  com- 
pound is  the  oxide,  Cu2O,  which  occurs  in  nature  as 
ruby  copper,  or  cuprite.  It  is  a  bright-red  substance  and 
can  easily  be  prepared  by  heating  copper  to  a  high  tem- 
perature in  a  limited  supply  of  air.  It  is  used  for  impart- 
ing a  ruby  color  to  glass.  By  treating  cuprous  oxide  with 
different  acids  a  number  of  cuprous  salts  can  be  made. 

Cuprous  chloride  (CuCl)  is  the  best-known  cuprous 
salt.  It  can  be  made  by  boiling  a  solution  of  cupric 
chloride  with  copper: 

CuCl2-f-Cu >-2CuCl 

It  is  a  white  salt,  nearly  insoluble  in  water. 

Cupric  compounds.  Cupric  salts  are  easily  made  by  dis- 
solving cupric  oxide  in  acids  or,  when  insoluble,  by  pre- 
cipitation. In  crystallized  form  most  of  them  are  blue  or 
green.  Since  they  are  so  much  more  familiar  than  the 
cuprous  salts,  they  are  frequently  called  merely  copper 
salts. 

Cupric  oxide  (CuO).  This  is  a  black  insoluble  sub- 
stance obtained  by  heating  copper  in  excess  of  air  or  by 
igniting  the  hydroxide  or  the  nitrate.  It  is  used  as  an 
oxidizing  agent. 

Cupric  sulfate  (CuSOJ.  When  crystallized  from  water, 
cupric  sulfate  forms  large  blue  crystals  of  the  hydrate 


COPPER,  MERCURY,  AND  SILVER  499 

CuSO4  •  5  H2O,  called  blue  vitriol,  or  bluestone.  The  salt  is 
a  by-product  in  silver  refining,  and  is  also  made  by  the 
oxidation  of  pyrite  containing  copper: 


The  salt  finds  extensive  use  in  electrotyping,  in  copper 
re  tin  ing,  as  a  remedy  for  hoof  diseases  (particularly  in 
sheep),  and  in  the  manufacture  of  insecticides.  Like  all 
copper  salts,  it  is  poisonous,  especially  to  lower  forms  of 
life.  When  added  even  in  very  minute  quantities  to 
water  containing  green  pond  scum  (algse),  the  plant  is 
quickly  killed.  Mixed  with  milk  of  lime  (which  precipi- 
tates copper  hydroxide)  it  constitutes  Bordeaux  mixture, 
which  is  used  as  a  spray  for  killing  molds  and  scale  on 
fruit  trees  and  vegetables. 

Cupric  sulfide  (CuS).  In  the  form  of  a  black  insoluble 
precipitate,  cupric  sulfide  (CuS)  is  easily  prepared  by  the 
action  of  hydrogen  sulfide  upon  a  solution  of  a  copper  salt : 

CuSO4  +  H2S  — >-  CuS  +  H2SO4 

It  is  insoluble  in  water  and  in  dilute  acids. 

Other  cupric  salts.  Among  the  other  cupric  salts  fre- 
quently used  in  the  laboratory  are  the  following,  most  of 
which  form  other  hydrates  in  addition  to  those  given : 

Cupric  nitrate  (Cu(NO3)2  •  6  H2O) :  blue  deliquescent  crystals 
Cupric  chloride  (CuCl2  •  2  H2O) :  light-blue  pearly  scales  or  needles 
Cupric  bromide  (CuBr2)  :  brownish-purple  crystals  resembling  iodine 
Cupric  acetate  (Cu(C2H3O2)2  •  H2O) :  a  blue,  easily  crystallized  salt 

Cupric  ammonia  compounds.  When  cupric  sulfate  in  solution 
is  treated  with  aqua  ammonia  the  insoluble  hydroxide  is  at 
first  precipitated : 

CuS04  +  2  NH4OH >-  Cu(OH)2  +  (NH4)2SO4 


500    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Continued  addition  of  ammonia  causes  the  precipitate  to  dis- 
solve, forming  an  intensely  colored  blue-purple  solution.  By 
adding  alcohol  to  the  solution  a  colored  salt  crystallizes, 
which  is  found  to  have  the  formula  Cu(NH8)4S04  •  H20.  In 
solution  this  salt  does  not  give  simple  cupric  ions  (Cu+  +),  but 
the  cation  is  Cu(NH8)4++  and  the  anion  is  S04".  All  soluble 
copper  salts  give  a  solution  of  this  same  blue-purple  color  when 
treated  with  an  excess  of  ammonia,  and  all  of  these  solutions 
contain  the  complex  cation  Cu(NH8)4+  +.  The  anion  may  be 
any  acid  radical,  so  that  we  have  such 
salts  as  the  following :  Cu(NH8)4Cl2 ; 
Cu(NH8)4(N08)2;  Cu(NH8)4C03.  Many 
other  metals  form  similar  salts, 
especially  cobalt,  nickel,  and  plati- 
num, and  these  are  sometimes  called 
ammine  salts. 

Electric  cells.  An  electric  cell 
is  a  device  for  converting  chem- 
ical energy  directly  into  electrical 
energy.  A  great  many  different 
chemical  reactions  can  be  arranged 
in  such  a  way  as  to  accomplish  this 
result,  and  the  combination  known 
as  the  Daniell  cell  will  serve  as 
an  illustration  of  the  most  familiar  types  of  cells.  In  this 
combination  two  plates,  one  of  copper  and  the  other  of 
zinc,  each  fashioned  so  as  to  have  a  large  surface,  are 
arranged  in  a  glass  jar,  as  shown  in  Fig.  195.  The  electro- 
lyte in  contact  with  the  zinc  plate  is  zinc  sulfate,  while 
that  in  contact  with  the  copper  plate  is  copper  sulfate. 

The  action  of  the  cell  can  be  explained  as  follows  :  The  zinc 
atoms  of  the  plate  A  have  a  tendency  to  leave  the  plate  and 
to  pass  into  solution  as  zinc  ions.  But  since  each  zinc  ion 
(Zn++)  carries  two  positive  charges,  the  formation  of  these 


FIG.  195.     A  simple   type 

of  electric  cell,  called  the 

Daniell  cell 


COPPER,  MERCURY,  AND  SILVER     501 

ions  leaves  the  plate  negatively  charged.  The  accumulation  of 
this  charge  soon  prevents  the  formation  of  more  zinc  ions  and 
the  process  stops  in  an  equilibrium. 

On  the  other  hand,  copper  ions  (Cu+  +)  from  the  dissolved 
copper  sulfate  have  a  tendency  to  deposit  upon  the  copper 
plate  B  as  metallic  atoms,  each  copper  ion  bringing  to  the 
plate  two  positive  charges.  The  accumulation  of  these  charges 
on  the  plate  soon  prevents  the  deposit  of  more  copper  ions, 
and  equilibrium  results. 

If  now  the  two  plates  are  joined  by  an  electrical  conductor 
(a  metal  wire)  the  negative  charge  on  the  zinc  plate  neutralizes 
the  positive  charge  on  the  copper  plate  by  flowing  through  the 
wire.  The  solution  of  the  zinc  and  the  deposit  of  the  copper 
can  now  proceed  as  long  as  the  two  plates  are  connected.  The 
chemical  reaction  taking  place  is  represented  by  the  equation : 

Zn  +  CuS04 >-  ZnS04  +  Cu  +  50,100  cal. 

In  this  particular  reaction  nearly  all  of  the  heat  is  converted 
into  electrical  energy,  which  may  be  obtained  from  the  wire 
and  may  be  in  turn  converted  into  light  or  work. 

The  order  of  the  metals  in  the  electromotive  series  (p.  191) 
is  the  order  of  intensity  with  which  the  metals  tend  to  pass 
into  ionic  form.  Any  two  metals  in  a  suitable  electrolyte  will 
constitute  a  cell  in  which  the  metal  higher  in  the  series  is  the 
negative  pole  and  the  lower  one  the  positive.  As  a  rule,  only 
a  part  of  the  chemical  energy  is  converted  into  electrical 
energy,  the  remainder  being  transformed  into  heat. 


MERCURY 

Occurrence.  Mercury  occurs  in  nature  chiefly  as  the 
sulfide  HgS,  called  cinnabar.  The  mercury  mines  of 
Spain  have  long  been  famous,  and  California  is  the 
next  largest  producer. 

Metallurgy.  Mercury  is  a  volatile  metal  which  has  but 
little  affinity  for  oxygen,  and  this  makes  the  metallurgy 


502    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

of  mercury  very  simple.  The  crushed  ore,  mixed  with 
a  small  amount  of  carbon  to  reduce  any  oxide  or  sulfate 
that  might  be  formed,  is  roasted  in  a  current  of  air.  The 
sulfur  burns  to  sulfur  dioxide,  while  the  mercury  vapor- 
izes and  is  condensed  in  a  series  of  condensing  vessels. 
The  metal  is  purified  by  distillation. 

Properties.  Mercury  is  a  heavy,  silvery  liquid,  with  a 
density  of  13.56.  It  boils  at  357°  and  solidifies  at  -38.9°. 
It  forms  alloys  (called  amalgams)  with  nearly  all  metals. 

Toward  acids  mercury  conducts  itself  very  much  like 
copper;  it  is  easily  attacked  by  nitric  acid  and  by  hot 
concentrated  su  If  uric  acid,  while  cold  sulfuric  acid  and 
hydrochloric  acid  have  no  effect  on  it. 

Uses.  Mercury  is  extensively  used'  in  the  construction 
of  many  scientific  instruments,  such  as  the  thermometer 
and  the  barometer,  and  as  a  liquid  over  which  to  collect 
gases  that  are  soluble  in  water.  The  readiness  with  which 
it  alloys  with  silver  and  gold  makes  it  very  useful  in  the 
extraction  o'f  these  elements.  All  salts  of  mercury  are 
made  directly  or  indirectly  from  the  purified  metal. 

Compounds  of  mercury.  Like  copper,  mercury  forms 
two  series  of  compounds :  the  mercurous  compounds,  of 
which  mercurous  chloride  (HgCl)  is  an  example;  and 
the  mercuric  compounds,  represented  by  mercuric  chloride 
(HgCl2). 

Mercuric  oxide  (HgO).  Mercuric  oxide  is  usually  obtained 
as  a  brick-red  substance  by  carefully  heating  the  nitrate : 

2  Hg(N03)2  — ^  2  HgO  +  4  NO2  +  O2 

It  can  also  be  obtained  in  a  yellow  form.  When  heated, 
the  oxide  darkens  until  it  becomes  almost  black ;  at  a 
higher  temperature  it  decomposes  into  mercury  and  oxy- 
gen  (p.  25). 


COPPER,  MERCURY,  AND  SILVER     503 

Mercuric  sulfate  (HgSOJ.  This  salt  is  easily  prepared 
by  the  action  of  an  excess  of  concentrated  sulfuric  acid 
upon  mercury,  the  reaction  being  like  that  of  sulfuric 
acid  upon  copper  (p.  247).  It  is  a  white  solid,  soluble 
in  dilute  sulfuric  acid,  but  hydrolyzed  into  a  yellow 
basic  salt  by  pure  water.  It  is  the  cheapest  salt  of 
mercury. 

Mercuric  chloride  (corrosive  sublimate)  (HgCl2).  This 
substance  can  be  made  by  dissolving  mercuric  oxide  in 
hydrochloric  acid.  On  a  commercial  scale  it  is  made  by 
heating  a  mixture  of  common  salt  and  mercuric  sulfate : 

2  NaCl  +  HgS04  — +  HgCl2  +  NaaSO4 

The  mercuric  chloride,  being  readily  volatile,  vaporizes 
and  is  condensed  again  in  cool  vessels.  It  is  a  white 
solid  and  is  soluble  in  water.  It  is  extremely  poisonous 
and  in  dilute  solutions  is  used  as  an  antiseptic  in  dressing 
wounds. 

Mercurous  chloride  (calomel)  (HgCl).  Being  insoluble, 
mercurous  chloride  is  precipitated  as  a  white  solid  when 
a  soluble  chloride  is  added  to  a  solution  of  mercurous 

HgN08  +  NaCl  — >-  HgCl  +  NaNO3 

Commercially  it  is  manufactured  by  heating  a  mixture  of 
mercuric  chloride  and  mercury.  It  is  a  common  medicine. 
Mercuric  sulfide  (HgS).  As  cinnabar,  this  substance 
forms  the  chief  native  compound  of  mercury,  and  occurs 
in  red  crystalline  masses.  By  passing  hydrogen  sulfide 
into  a  solution  of  a  mercuric  salt,  mercuric  sulfide  is 
precipitated  as  a  black  powder  insoluble  in  water  and 
acids.  By  other  means  it  can  be  prepared  as  a  brilliant 
red  powder,  known  as  vermilion,  which  is  used  as  a 
pigment  in  fine  paints. 


504    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Other  salts  of  mercury.  The  following  salts  have  fre- 
quent uses  in  the  laboratory : 

Mercurous  nitrate  (HgNO3  •  2  H2O)  :  colorless  crystals  soluble  in 
dilute  nitric  acid 

Mercuric  nitrate,  Hg(NO3)2  •  8  H2O :  colorless  crystals  soluble  in 
dilute  nitric  acid 

Mercurous  sulfate  (Hg2SO4)  :  white  crystalline  solid  nearly  insolu- 
ble in  dilute  acids 

Mercuric  fulminate,  Hg(ONC)2 :  a  white  solid  made  by  the  action 
of  nitric  acid  upon  mercury  in  the  presence  of  alcohol.  It  is 
used  as  the  explosive  material  in  percussion  primers 

SILVER 

Occurrence.  Silver  is  found  in  small  quantities  in  the 
uncombined  state ;  usually,  however,  it  occurs  in  com- 
bination with  sulfur,  either  as  the  sulfide  Ag2S  or  as  a 
constituent  of  other  sulfides,  especially  those  of  lead, 
copper,  antimony,  and  arsenic.  It  is  also  found  alloyed 
with  gold. 

In  this  country  silver  is  produced  almost  entirely  in 
connection  with  lead,  and  it  will  be  convenient  to  con- 
sider the  metallurgy  of  the  two  metals  together  in  the 
next  chapter. 

The  United  States  produces  about  one  third  of  the 
world's  output  of  silver,  and  America,  including  Mexico 
and  Canada,  about  70  per  cent. 

The  refining  of  silver.  Crude  silver  obtained  by  any 
process  may  contain  a  number  of  metals,  especially  copper 
and  gold,  and  is  usually  refined  either  by  parting  with 
sulfuric  acid  or  by  electrolysis. 

I.  Cupellation  and  parting  with  sulfuric  acid.  In  this  proc- 
ess the  impure  metal  is  heated  on  an  open  hearth  in  a 
strong  current  of  air.  The  various  metallic  impurities 


COPPER,  MERC  UK  Y,  AND  SILVER  505 

(excepting  gold)  are  in  this  way  largely  converted  into 
oxides  and  swept  off  as  dross,  leaving  the  silver  alloyed 
with  small  percentages  of  gold,  copper,  and  iron.  It  is 
then  cast  into  ingots  known  as  dorS  bars,  since  they 
contain  gold. 

In  order  to  recover  the  gold  the  alloy  is  treated  with 
hot  concentrated  sulfuric  acid,  which  converts  into  su-lfates 
all  the  metals  except  the  gold.  When  water  is  added  to 
the  resulting  mixture  the  sulfates  of  copper,  silver,  and 
iron  pass  into  solution,  while  the  gold,  together  with  the 
lead  sulfate  and  any  unattacked  substances,  settles  as 
a  mud  from  which  the  gold  is  subsequently  recovered. 
The  silver  is  separated  from  the  solution  of  the  sulfates 
by  suspending  in  the  latter  clean  copper  plates,  the  copper 
displacing  the  silver,  which  is  deposited  in  crystalline  form : 

AgaS04  4-  Cu  — >•  CuSO4  +  2  Ag 

The  copper  sulfate  obtained  as  a  by-product  in  this  process 
furnishes  much  of  the  blue  vitriol  of  commerce. 

2.  Electrolytic  refining.  Electrolysis  of  the  impure  silver 
is  now  carried  out  extensively,  the  process  being  con- 
ducted in  a  way  very  similar  to  the  electrolysis  of 
copper.  The  electrolyte  used  is  a  solution  of  silver 
nitrate  in  nitric  acid.  The  silver  is  deposited  as  crystals, 
which  are  mechanically  brushed  off  the  cathode,  collected, 
and  melted  into  bars. 

Properties  of  silver.  Silver  is  a  heavy,  rather  soft,  white 
metal,  very  ductile  and  malleable,  and  capable  of  taking 
a  high  polish.  It  surpasses  all  other  metals  as  a  con- 
ductor of  heat  and  electricity,  but  is  too  costly  to  find 
extensive  use  for  such  purposes.  It  melts  at  a  little  lower 
temperature  than  copper.  It  alloys  readily  with  other 
heavy  metals,  and  when  it  is  to  be  used  for  coinage  or 


506    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Fiu.  196.    The  process  of  silver  plating 


for  tableware,  a  small  amount  of  copper  —  from  8  per 
cent  to  10  per  cent  — is  melted  with  it  to  give  it  hard- 
ness (sterling  silver). 

It  is  not  acted  upon  by  water  or  air,  but  is  quickly 
tarnished  when  in  contact  with  sulfur  compounds  (eggs, 
mustard,  perspiration),  turning  quite  black  in  time.  Hy- 
drochloric acid  and  fused  alkalies  do  not  act  upon  it,  but 
nitric  acid  and  hot  concentrated  sulfuric  acid  dissolve  it 
with  ease.  When  a  solution  of  a  silver  salt  is  treated 

with  a  strong 
reducing  agent, 
metallic  silver 
is  precipitated. 
Under  proper 
conditions  this 
takes  the  form 
of  a  brilliant 

mirror  deposited  on  the  sides  of  the  glass  vessel.  Mirrors 
and  glass  reflectors  are  usually  made  in  this  way. 

Electroplating  with  silver.  Since  silver  is  not  acted  upon  by 
water  or  air  and  has  a  pleasing  appearance,  it  is  used  to  coat 
various  articles  made  of  cheaper  metals.  Such  articles  are  said 
to  be  silver  plated,  and  the  process  by  which  this  is  done  is  very 
similar  to  electroplating  with  nickel  (p.  490).  The  object  to 
be  plated  (as,  for  example,  a  spoon)  is  attached  to  a  wire  and 
dipped  into  a  solution  of  a  suitable  silver  salt.  Electrical  con- 
nection is  made  in  such  a  way  that  the  article  to  be  plated  k 
the  cathode  (Fig.  196),  while  the  anode  A  is  made  up  of  one 
or  more  plates  of  silver. 

Compounds  of  silver.  Silver  forms  only  one  series  of  salts, 
which  corresponds  to  the  mercurous  and  the  cuprous  series. 

Silver  nitrate  (lunar  caustic)  (AgN03).  This  salt  is  easily 
prepared  by  dissolving  silver  in  nitric  acid  and  evaporating 


COPPER,  MERCURY,  AND  SILVER  507 

the  resulting  solution.  It  crystallizes  in  flat,  colorless 
plates,  and  when  heated  carefully  can  be  melted  without 
decomposition.  When  cast  into  sticks  it  is  called  lunar 
caustic,  for  it  has  a  very  corrosive  action  on  flesh  and  is 
sometimes  used  in  surgery  to  burn  away  abnormal  growths. 

The  alchemists  designated  the  metals  by  the  names  of  the 
heavenly  bodies.  The  moon  (luna)  was  the  symbol  for  silver ; 
hence  the  name  lunar  caustic. 

Some  less  important  salts.  The  following  is  a  list  m  of 
some  of  the  silver  salts  used  in  the  arts  or  in  analysis: 

Silver  sulfate  (Ag2SO4)  :  a  sparingly  soluble  salt  crystallizing  in 
white  needles 

Silver  nitrite  (AgNO2)  :  a  white  salt  soluble  in  hot  water 

Silver  sulfide  (Ag2S)  :  found  in  nature  as  argentite.  It  is  insoluble 
both  in  water  and  in  acids  and  may  be  prepared  by  precipita- 
tion as  a  black  powder 

Silver  cyanide  (AgCN) :  a  curdy  white  precipitate  insoluble  in 
water 

Silver  chromate  (Ag2CrO4) :  a  brick-red  solid  prepared  by  pre- 
cipitation 

Compounds  of  silver  with  the  halogens.  The  chloride 
(AgCl),  the  bromide  (AgBr),  and  the  iodide  (Agl)  are 
insoluble  in  water  and  in  acids  and  therefore  are  precipi- 
tated by  bringing  together  a  soluble  halogen  salt  with 
silver  nitrate : 

AgNO8  +  KC1  — >•  AgCl  +  KNO, 

These  salts  are  remarkable  for  the  fact  that  they  are  very 
sensitive  to  the  action  of  light,  undergoing  a  change  of 
color  and  chemical  composition  when  exposed  to  sunlight. 
It  is  thought  that  this  change  consists  in  a  partial  decom- 
position according  to  the  equilibrium  equation : 


508    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


FIG.  197.   The  negative  plate 


If  some  material  is  present  that  will  absorb  chlorine,  such 
as  gelatin,  the  action  is  more  pronounced.  In  this  case 
the  gelatin  is  called  a  sensitizer.  It  is  upon  these  general 

facts    that    the    art    of 
photography   is   based. 

Photography.  From  a 
chemical  standpoint  the 
processes  of  photography 
may  be  described  under 
two  heads  :  (1)  the  prepa- 
ration of  the  negative ; 
(2)  the  preparation  of 
the  print. 

1.  Preparation  of  the  negative.  The  plate  used  in  the  prepara- 
tion of  the  negative  is  made  by  spreading  a  thin  layer  of  gela- 
tin, in  which  colloidal  silver  bromide  is  suspended  (silver 
iodide  is  sometimes  added  also),  over  a  glass  plate  or  celluloid 
film  and  allowing  it  to  dry.  When  the  plate  so  prepared  is 
placed  in  a  camera  and  the  image  of  some  object  is  focused 
upon  it,  the  silver  salt  undergoes  a  change  which  is  propor- 
tional at  each  point  to  the  intensity  of  the  light  falling  upon  it. 
In  this  way  an  image  of 
the  object  photographed 
is  produced  upon  the 
plate.  This  image,  how- 
ever, is  invisible  and  is 
therefore  called  latent.  It 
can  be  made  visible  by 
the  process  of  developing. 
To  develop  the  image 
the  exposed  plate  is  im-  FIG.  198.  The  positive  print 

mersed  in    a  solution  of 

some  reducing  agent  called  the  developer.  While  the  developer 
will  in  time  reduce  all  the  silver  salt,  it  acts  much  more 
rapidly  upon  that  which  has  been  exposed  to  the  light.  The 
action  is  therefore  continued  only  long  enough  to  bring  out 


COPPER,  MERCURY,  AND  SILVER     509 

the  image.  The  reduced  silver  is  deposited  in  the  form  of 
a  black  film  which  adheres  closely  to  the  plate. 

The  unaffected  silver  salt  is  now  removed  from  the  plate  by 
immersing  it  in  a  solution  of  sodium  thiosulfate  (hypo).  The 
plate  is  then  washed  with  water  and  dried.  The  plate  so  pre- 
pared is  called  the  negative  (Fig.  197),  because'  it  is  a  picture 
of  the  object  photographed,  with  the  lights  and  positions 
exactly  reversed. 

2.  Preparation  of  the  print.  The  print  is  made  on  paper  which 
is  prepared  in  much  the  same  way  as  the  negative  plate.  The 
negative  is  placed  upon  this  paper  and  exposed  to  the  light  in 
such  a  way  that  the  light  must  pass  through  the  negative  before 
striking  the  paper.  If  the  paper  is  coated  with  silver  chloride, 
a  visible  image  is  produced,  in  which  case  a  developer  is  not 
needed.  Proofs  are  made  in  this  way.  In  order  to  make  them 
permanent  the  unchanged  silver  chloride  must  be  dissolved 
off  with  sodium  thiosulfate.  The  print  is  then  toned  by  dip- 
ping it  into  a  solution  of  gold  or  platinum  salts,  in  which 
process  the  silver  on  the  print  passes  into  solution,  while  the 
gold  or  platinum  takes  its  place.  These  metals  give  a  charac- 
teristic color  or  tone  to  the  print,  the  gold  making  it  reddish 
brown,  while  the  platinum  gives  it  a  steel-gray  tone.  Since 
the  darkest  places  on  the  negative  cut  off  the  most  light,  it 
is  evident  that  the  lights  of  the  print  (Fig.  198)  will  be  the 
reverse  of  those  of  the  negative,  and  will  therefore  correspond 
to  those  of  the  object  photographed. 

EXERCISES 

1.  Why  has  copper  or  bronze  been  used  for  so  long  a  time? 

2.  Why  do  we  have  so  many  relics  from  the  bronze  age  and  so 
few  from  the  iron  age  ? 

3.  Why  is  a  solution  of  copper  sulfate  acid  to  litmus  paper? 

4.  How  should  you  account  for  the  fact  that  so  many  different 
salts  of  copper  have  the  same  blue  color  in  dilute  solutions  ? 

5.  How  could  you  distinguish  between  mercurous  chloride  and 
mercuric  chloride? 


510    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

6.  Crude  silver  usually  contains  iron  and  lead.    What  would 
become  of  these  in  refining  silver  by  parting  with  sulf uric  acid  ? 

7.  Mercuric  nitrate  and  silver  nitrate  are  both  white  soluble 
solids.    How  could  you  distinguish  between  them? 

8.  How  do  you  account  for  the  fact  that  a  silver  spoon  gradu- 
ally darkens  when  in  contact  with  eggs? 

9.  Why  are  copper,  mercury,  and  silver  all  found  to  some  extent 
native  in  nature  ? 

10.  What  properties  make  mercury  useful  in  thermometers  and 
in  barometers  ? 

11.  Write  equations  for  the  action  of  concentrated  sulf  uric  acid 
on  mercury ;  of  nitric  acid  or.  silver. 

12.  Suggest  a  method  for  obtaining  pure  silver  from  a  silver  coin. 

13.  When  a  solution  of  silver  nitrate  is  added  to  a  solution  of 
potassium  chlorate,  no  precipitate  forms.    How  do  you  account  for 
the  fact  that  a  precipitate  of  silver  chloride  is  not  formed  ? 

14.  Which  of  the  ores  of  copper  contains  the  largest  percentage 
of  copper? 

15.  What  would  be  the  effect  of  immersing  an  unexposed  photo- 
graphic film  in  a  solution  of  sodium  thiosulfate? 

16.  One  ton  of  cinnabar  will  yield  how  many  pounds  of  mercury  ? 

17.  Suppose  you  wish  to  prepare  1000  g.  of  calomel,  what  sub- 
stances and  what  weight  of  each  would  be  required? 

18.  A  silver  dollar  weighs  approximately  26.5  g.    What  weight  of 
silver  nitrate  could  be  prepared  from  such  a  coin? 


CHAPTER  XL 
TIN  AND  LEAD 


NAME 

SYMBOL 

ATOMIC 

WEIGHT 

DENSITY 

MELTING 

POINT 

COMMON  OXIDES 

Tin    .... 

Lead      .     .     . 

Sn 
Pb 

118.7 
207.20 

7.3 
11.37 

231.9° 

327.4° 

SnO                SnO2 
PbO  Pb804   Pb02 

TIN 

Occurrence.  Tin  is  found  in  nature  chiefly  as  the  oxide 
SnO2,  called  cassiterite,  or  tinstone.  The  most  famous  mines 
are  in  Cornwall,  in  the  East  Indies,  and  in  Bolivia.  Very 
little  is  produced  in  the  United  States  from  domestic  ores. 

Metallurgy.  In  principle  the  metallurgy  of  tin  is  very 
simple.  The  ore,  separated  as  far  as  possible  from  earthy 
materials,  is  mixed  with  carbon  and  heated  in  a  furnace, 
the  reduction  taking  place  readily.  The  equation  is 

SnO2-f  C ^Sii  +  CO2 

The  chief  difficulty  is  in  freeing  the  tin  from  iron,  and 
commercial  tin  often  contains  a  considerable  percentage  of 
the  latter  element.  The  metal  is  often  purified  by  care- 
fully heating  it  until  it  is  partly  melted;  the  pure  tin 
melts  first  and  can  be  drained  away  from  the  impurities. 
Properties.  Pure  tin,  called  block  tin,  is  a  soft  white 
metal  with  a  silverlike  appearance  and  luster;  it  melts 
readily  and  is  somewhat  lighter  than  copper,  having  a 
density  of  7.3.  It  is  malleable  and  can  be  rolled  out  into 

511 


512    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

very  thin  sheets,  forming  tin  foil  ;  most  tin  foil,  however, 
contains  a  considerable  percentage  of  lead. 

Under  ordinary  conditions  tin  is  unchanged  by  air  or 
moisture,  but  at  a  high  temperature  it  burns,  forming  the 
oxide  SnO2.  Dilute  acids  have  little  effect  upon  it,  but 
concentrated  acids  attack  it  readily.  Concentrated  hydro- 
chloric acid  changes  it  into  the  chloride  : 
Sn  +  2HCl 


With  sulfuric  acid,  stannous  sulfate  and  sulfur  dioxide 
are  formed  : 

Sn  +  2  H2SO4  -  >•  SnSO4  +  SO2  +  2  H2O 

Concentrated  nitric  acid  oxidizes  it,  forming  a  white  in- 
soluble compound,  H2SnO3,  called  metastannic  acid. 

Uses  of  tin.  A  great  deal  of  tin  is  used  in  the  making 
of  tin  plate.  The  process  consists  in  dipping  thin  sheets  of 
iron  into  the  melted  tin  and  is  quite  similar  to  that  of 
galvanizing  iron  (p.  447).  Owing  to  its  resistance  to  the 
action  of  air  and  weak  acids,  tin  plate  is  used  in  many 
ways,  such  as  in  roofing  and  in  the  manufacture  of  tin 
cans,  cooking  vessels,  and  similar  articles.  Small  pipes  of 
block  tin  are  used  instead  of  lead  for  conveying  pure 
water  or  liquids  containing  dilute  acids,  such  as  soda 
water.  Many  useful  alloys  contain  tin  (p.  497).  Pewter 
and  soft  solder  are  alloys  of  tin  and  lead. 

Rusting  of  tin  plate.  If  the  coating  of  tin  on  tin  plate 
is  scratched  through  to  the  iron,  the  iron  will  rust  faster 
than  if  there  were  no  tin  covering.  The  two  metals  and 
the  water  constitute  a  battery,  much  like  the  Daniell  cell 
(p.  500),  and  in  a  battery  the  metal  highest  in  the  dis- 
placement series  is  the  one  that  is  corroded.  In  the  case 
of  galvanized  iron  the  zinc  rusts  first,  and  the  iron  resists 
rusting  as  long  as  any  zinc  is  present. 


TIN  AND  LEAD  513 

Soldering  and  brazing.  The  use  of  solder  in  joining  two  metal 
surfaces  depends  upon  (1)  the  low  melting  point  of  the  solder 
and  (2)  the  fact  that  it  flows  over  clean  metal  surfaces  and 
sticks  to  them  on  cooling.  To  secure  clean  surfaces  free  from 
oxide  a  suitable  flux  must  be  used  which  will  either  dissolve 
the  oxide  as  fast  as  it  forms  or  will  reduce  it  again  to  metal. 
The  usual  fluxes  are  zinc  chloride,  ammonium  chloride,  rosin, 
and  stearin.  In  brazing,  or  hard  soldering,  the  process  is  essen- 
tially the  same,  except  that  a  low-melting  brass  is  used  instead 
of  solder,  and  borax  is  used  as  a  flux. 

Compounds  of  tin.  Tin  forms  two  series  of  compounds : 
the  stannous,  in  which  the  tin  is  bivalent,  illustrated  in  the 
compounds  SnO,  SnS,  SnCl2 ;  the  stannic,  in  which  it  is 
tetravalent,  as  shown  in  the  compounds  SnO2,  SnS2.  There 
is  also  an  acid,  H  SnO  ,  called  stannic  acid,  which  forms 

2  37 

a  series  of  salts  called  stannates.    Only  a  few  compounds 
of  tin  need  be  described. 

Stannic  oxide  (Sn02).  Stannic  oxide  is  of  interest,  since 
it  is  the  chief  compound  of  tin  found  in  nature.  It  is 
sometimes  found  in  good-sized  crystals,  but  as  prepared  in 
the  laboratory  it  is  a  white  powder.  When  stannic  oxide 
is  fused  with  potassium  hydroxide  it  forms  potassium 
stannate,  acting  very  much  like  silicon  dioxide: 

SnO2  +  2  KOH >•  K2SnO8  +  H2O 

Stannic  acid  (H2Sn03).  The  compound  H2SnO8  can  be 
prepared  in  a  number  of  ways,  and  it  is  a  white  solid, 
insoluble  in  water.  It  is  a  remarkable  fact  that  its  clirm- 
ical  properties  depend  to  a  large  extent  upon  the  way  in 
which  it  is  made.  Some  methods  yield  a  product  that  is 
readily  soluble  in  dilute  acids,  while  the  product  formed 
by  the  action  of  concentrated  nitric  acid  upon  tin  is  insol- 
uble in  acids  and  is  called  metastannic  acid.  The  soluble 
forms  are  colloidal  in  character,  and  it  is  probable  that 


514    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

the  differences  in  the  precipitated  colloids  are  due  to  the 
size  of  the  particles  formed  rather  than  to  a  different 
chemical  structure  of  the  compounds. 

Chlorides  of  tin.  Stannous  chloride  is  prepared  by  dis- 
solving tin  in  concentrated  hydrochloric  acid  and  evapo- 
rating the  solution  to  crystallization.  The  crystals  which 
are  obtained  have  the  composition  SnCl2  •  2  H2O  and  are 
known  as  tin  crystals.  By  treating  a  solution  of  stannous 
chloride  with  aqua  regia,  stannic  chloride  is  formed : 

SnCl2  +  2  [Cl] >-  SnCl4 

The  salt  which  crystallizes  from  such  a  solution  has  the 
composition  SnCl4  •  5  H2O  and  is  known  commercially  as 
oxymuriate  of  tin.  If  metallic  tin  is  heated  in  a  current  of 
dry  chlorine,  the  anhydrous  chloride  (SnCl4)  is  obtained 
as  a  heavy  colorless  liquid  which  fumes  strongly  on 
exposure  to  air.  A  great  deal  of  stannic  chloride  is  re- 
covered from  scrap  tin  plate  by  detinning  with  chlorine. 

The  ease  with  which  stannous  chloride  takes  up  chlorine 
to  form  stannic  chloride  makes  it  a  good  reducing  agent  in 
many  reactions,  changing  the  higher  chlorides  of  metals  to 
lower  ones.  Thus  mercuric  chloride  is  changed  into  mercurous 

chloride : 

SnCl2  +  2  HgCl2 >-  SnCl4  +  2  HgCl 

If  the  stannous  chloride  is  in  excess  the  reaction  may  go 
further,  producing  metallic  mercury : 

SnCl2  +  2  HgCl >•  SnCl4  +  2  Hg 

Ferric  chloride  is  in  like  manner  reduced  to  ferrous  chloride : 
SnCl2  +  2  FeCl8 >-  SnCl,  +  2  FeCl, 

The  chlorides  of  tin,  as  well  as  the  alkali  stannates,  are 
much  used  as  mordants  in  dyeing  processes.  The  hydroxides 
of  tin  and  free  stannic  acid,  which  are  easily  liberated  from 


TIN  AND  LEAD  515 

these  compounds,  possess  in  very  marked  degree  the  power  of 
fixing  dyes  upon  fibers,  as  explained  under  aluminium  (p.  458). 

Other  compounds  of  tin.  A  few  of  the  other  important 
compounds  are  included  in  the  following  list: 

Stannous  sulfide  (SnS)  :  a  nearly  black  precipitate  formed  by  the 
action  of  hydrogen  sulfide  upon  a  stannous  salt 

Stannic  sulfide  (SnS2)  :  a  bright-yellow  precipitate  formed  by  the 
action  of  hydrogen  sulfide  upon  a  stannic  salt 

Ammonium  chlorostannate,  (NH4)2SnCl6:  a  well-crystallized  salt 
used  as  a  mordant  in  dyeing,  and  called  pink  salt 

LEAD 

Occurrence.  Lead  is  found  in  nature  chiefly  as  the 
sulfide  PbS,  called  galenite.  In  the  United  States  this  is 
mined  principally  in  Missouri  (Fig.  199)  and  Idaho. 

Metallurgy  of  lead.  Almost  all  the  lead  of  commerce 
is  made  from  galenite,  which  usually  contains  some  silver 
To  obtain  this  silver  most  economically  it  is  customary 
to  combine  richer  silver  ores  with  lead  ores  and  work  the 
two  together. 

Reduction  of  silver-bearing  lead  ores.  The  sulfide  ores 
are  first  roasted  until  a  part  of  the  sulfide  has  been 
changed  into  the  oxide  and  the  sulfate.  The  air  is  then 
shut  off  and  the  heating  continued,  which  brings  about 
the  reactions  indicated  in  the  following  equations: 

2  PbO  +  PbS  — *•  3  Pb  4-  SO2 
PbSO4  +  PbS  — )-  2  Pb  +  2  S02 

By  reactions  which  are  similar  to  the  above  the  silver 
compounds  are  reduced,  the  silver  alloying  with  the  lead.. 
In  some  more  recent  plants  the  ores  are  roasted  until 
nearly  all  of  the  sulfide  has  been  oxidized.  The  roasted 


516    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

ore  is  then  mixed  with  coke  and  heated  in  a  furnace  like 
a  copper  matte  furnace.  All  lead  oxide  is  reduced  to  lead, 
while  the  lead  sulfate  is  reduced  to  lead  sulfide  and  forms 
a  liquid  matte.  This  is  drawn  off  and  once  more  roasted. 
The  softening  of  lead.  The  lead  obtained  in  this  way  is 
called  hard  lead,  and  in  addition  to  silver  contains  smaller 
quantities  of  other  elements,  especially  of  copper,  arsenic, 


FIG.  199.   Mining  galenite  in  the  Joplin  region  in  Missouri 

antimony,  gold,  and  bismuth.  The  lead  is  softened  by  melt- 
ing it  in  an  open  furnace  with  free  access  of  air.  This 
converts  most  of  the  impurities  (as  well  as  some  lead)  into 
oxides  which  float  upon  the  melted  lead  and  can  easily  be 
removed.  The  partially  purified  lead  is  called  soft  lead. 

Desilverizing  of  lead  (the  Parkes  process).  The  lead  is  melted 
in  large  kettles  holding  as  much  as  30  tons,  and  about  1  per  cent 
of  zinc  is  stirred  in.  These  two  metals  do  not  mix  to  any  great 
extent,  and  gold,  silver,  and  copper  are  much  more  soluble  in 
zinc  than  in  lead.  Consequently  when  the  stirring  ceases,  the 
zinc  rises  to  the  surface  of  the  lead,  carrying  with  it  the  other 


TIN  AND  LEAD  51 T 

metals.  The  zinc  is  then  skimmed  off  and  distilled  (by  which 
process  the  zinc  is  recovered  to  be  used  again),  and  the  residue 
of  silver  and  gold  is  melted  down  and  cast  into  ingots  called 
dore  bars.  These  are  refined  as  explained  under  silver  (p.  505). 
An  electrolytic  method  (Betts  process]  is  now  being  used 
similar  to  the  one  employed  with  copper,  but  with  many 
special  details. 

Properties.  Lead  is  a  heavy  metal  which  has  a  brilliant 
silvery  luster  on  a  freshly  cut  surface,  but  which  soon 
tarnishes  by  oxidation  to  a  dull  blue-gray  color.  It  is  soft, 
easily  fused  (melting  at  327.4°),  and  malleable,  but  has 
little  toughness  or  strength. 

It  is  not  acted  upon  to  any  great  extent,  under  ordinary 
conditions,  by  the  oxygen  of  the  air,  but  at  a  high  tem- 
perature is  changed  into  the  oxide.  .With  the  exception 
of  hydrochloric  and  sulfuric  acids  (which  form  insoluble 
compounds),  most  acids,  even  very  weak  ones,  act  upon  it, 
forming  soluble  lead  salts.  Hot  concentrated  hydrochloric 
and  sulfuric  acids  also  attack  it  to  a  slight  extent. 

Uses.  Lead  finds  many  important  applications  in  the 
industries,  chiefly  in  the  manufacture  of  storage  batteries, 
hi  linings  for  sulfuric  acid  plants,  in  alloys  of  various  kinds 
(such  as  shot,  antifriction  metals,  type  metal,  and  pewter), 
and  in  water  pipes  for  plumbing.  Since  lead  dissolves  to 
some  extent  in  pure  water,  it  should  not  be  used  for 
pipes  that  are  to  carry  rain  water.  About  one  third  of 
the  annual  production  of  lead  is  used  in  making  paint, 
and  is  permanently  lost. 

Compounds  of  lead.  In  nearly  all  its  compounds  lead  is 
bivalent,  but  in  a  few  of  its  compounds  it  has  a  valence 
of  four.  All  of  its  compounds  are  poisonous. 

Lead  oxides.  Lead  forms  a  number  of  oxides,  the  most 
important  of  which  are  the  following: 


518    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

1.  Litharge,  massicot  (PbO).    This  oxide  forms  when  hot 
lead  is  oxidized  by  the  air.     If  the  temperature  is  high 
enough  to  melt  the  oxide,  the   color  of  the   product  is 
deep  yellow  or  even  red,  and  the  oxide  is  called  litharge. 
It  is  used  in  the  preparation  of  boiled  linseed  oil  and  in 
glassmaking.    If  the  oxide  is  not  melted,  it  forms  a  light 
yellow  powder  called  massicot,  which  is  used  in  making 
red  lead. 

2.  Red  lead,  or  minium  (Pb3OJ.    Minium  is  prepared  by 
heating  lead  (or  massicot)  to  a  high  temperature  in  contact 
with  a  current  of  air.    It  is  a  heavy  powder  of  a  beautiful 
red   color  and  is  much  used  as  a  pigment  for  painting 
structural  iron.    Mixed  with  linseed  oil  it  forms  a  cement 
used  in  joining  gas  pipes. 

3.  Lead  peroxide  (Pb02).    This  is  left  as  a  residue  when 
minium  is  heated  with  nitric  acid : 

PbgO4  +  4  HNO8  — **  2  Pb(N03)2  +  PbO2  +  2  H2O 

It  is  a  brown  powder  which  easily  gives  up  a  part  of  its 
oxygen  and,  like  manganese  dioxide  and  barium  dioxide, 
is  a  good  oxidizing  agent. 

Lead  sulfide  (PbS).  In  nature  this  compound  occurs  in 
a  highly  crystalline  form  called  galenite,  the  crystals  having 
much  the  same  color  and  luster  as  pure  lead.  It  is  readily 
prepared  in  the  laboratory  as  a  black  precipitate,  by  the 
action  of  hydrogen  sulfide  upon  soluble  lead  salts : 

Pb(N03)2  +  H2S  — +  PbS  +  2  HN08 

It  is  insoluble  both  in  water  and  in  dilute  acids. 

Lead  carbonate.  While  the  normal  carbonate  of  lead, 
PbCO3,  is  found  to  some  extent  in  nature  and  can  be  pre- 
pared in  the  laboratory,  basic  carbonates  of  varying  com- 
position are  much  more  easy  to  obtain.  One  of  the  simplest 


TIN  AND  LEAD 


519 


FIG.  200.   A  crock  filled 

with  thin  lead  plates  for 

making  white  lead 


of  these  has  the  composition  2  PbCO8  .  Pb(OH)2  and  is 
called  white  lead.  This  is  prepared  on  a  large  scale  as  a 
white  pigment  and  as  a  body  for 
paints  which  are  to  be  colored  with 
other  substances. 

Manufacture  of  white  lead.  White 
lead  can  be  prepared  by  a  number  of 
processes,  but  no  other  seems  to  pro- 
duce a  product  of  as  desirable  physical 
properties  as  the  old  Dutch  process, 
which  has  been  used  for  centuries, 
though  with  many  improvements.  In 
this  process  the  lead  is  cast  into  per- 
forated plates  called  buckles,  which  are 
placed  loosely  upon  each  other  in  a 
crock  of  the  shape  shown  in  Fig.  200, 
the  ledge  B,  formed  by  the  constriction  of  the  crock,  support- 
ing the  plates.  Under  them  in  .4  is  poured  a  suitable  quantity 
of  dilute  acetic  acid,  and  the  crocks  so  charged  are  placed 
in  banks  and  covered  with  stable  manure  or  spent  tanbark. 
The  heat  of  fermentation  in  the  latter  warms  the  acid,  the 
fumes  of  which  at- 
tack the  lead,  form- 
ing lead  acetate.  The 
carbon  dioxide  from 
the  fermentation  en- 
ters into  reaction  with 
the  acetate,  producing 
the  basic  carbonate 
and  regenerating  acetic 
acid.  This  acid  acts  FlG  201  Lead  buckleg  before  and  after 
again  upon  the  lead,  exposure  to  acetic  acid  and  carbon  dioxide 
and  the  process  contin- 
ues until  the  buckles  are  almost  completely  converted  into 
the  desired  compound.  Fig.  201  shows  a  buckle  before  and 
after  the  corrosion,  the  white  carbonate  resembling  enamel. 


520    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

A  more  modern  process,  known  as  Carter's*  process,  is  much 
used.  Very  finely  powdered  lead  is  placed  in  a  rotating  barrel, 
and  dilute  acetic  acid  is  sprayed  in.  Carbon  dioxide  is  admitted 
at  the  same  time,  and  the  process  is  continued  for  about  two 
weeks. 

Other  important  compounds  of  lead.  Soluble  salts  of  lead 
are  obtained  by  dissolving  litharge  in  the  appropriate  acid. 
Insoluble  salts  are  obtained  by  precipitation.  Some  of  the 
most  important  are  the  following: 

Lead  nitrate,  Pb(NO3)2 :  white  soluble  crystals 
Lead  chloride  (PbCl2) :  white  needles,  very  sparingly  soluble 
Lead  sulfate  (PbSO4) :  an  insoluble  white  crystalline  powder 
Lead   acetate,   Pb(C2H3O2)2  •  3  H2O:    a   soluble   white   salt  called 

sugar  of  lead 
Lead  chromate  (PbCrO4) :  used  as  a  pigment  in  paint  (chrome  yellow) 

Paints.  A  paint  consists  of  three  essential  ingredients: 
the  vehicle,  the  body,  and  the  pigment. 

1.  The  vehicle,  or  liquid  medium.    This  must  be  an  oil  which 
will  dry  rapidly  and  harden  in  drying  to  a  more  or  less  flexible, 
hornlike  body.    These  changes  in  the  oil  are  due  to  oxidation 
by  the  air.    A  number  of  different  oils  will  serve  this  purpose, 
but  linseed  oil  has  long  been  used  as  the  standard  drying  oil, 
since  it  can  be  produced  in  quantity  and  at  moderate  cost.    It 
is  customary  to  add  to  it  a  dryer,  made  by  boiling  some  of  the 
oil  with  oxides   of  manganese,  lead,  or   cobalt.    The  oxides 
enter  into  combination  with  the  oil  and  assist  catalytically  in 
its  oxidation. 

2.  The  body.    The  body  of  the  paint  must  be  some  solid  mate- 
rial, suspended  in  the  oil,  which  will  give  a  smooth  and  waxy 
surface  as  the  paint  dries,  and  will  have  good  covering  power. 
While  white  lead  meets  these  requirements,  it  is  moderately 
expensive  and  it  also  blackens  when  exposed  to  hydrogen  sul- 
fide,  which  is  likely  to  be  present  in  the  air  in  cities.    Other 
bodies  are  now  frequently  combined  with  the  lead,  or  replace  it 
altogether,  among  them  being  zinc  oxide,  China  clay  (or  kaolin), 


TIN  AND  LEAD 


521 


barium  sulfate,  and  a  product  called  lithopone  (p.  448).  For 
some  purposes  these  materials  are  a  real  advantage,  and  they 
are  not  to  be  regarded  as  adulterants  unless  sold  as  white  lead. 

3.  The  pigment,  or  coloring  matter.  In  the  case  of  white  paints 
the  body  serves  also  as  the  coloring  matter.  For  other  colors 
a  specific  pigment  must  be  added.  In  most  cases  these  pig- 
ments are  metallic  oxides  or  salts  and  are  frequently  natural 
products.  Sometimes 
they  are  prepared  by 
precipitating  an  amor- 
phous body  (usually  a 
colloid)  in  the  presence 
of  an  organic  dye,  the 
dye  being  absorbed  by 
the  precipitate  and  giv- 
ing it  a  color.  Such 
pigments  can  be  pre- 
pared in  an  endless 
variety  of  colors  and 
are  called  lakes.  They 
are  usually  not  so  per- 
manent as  mineral  pig- 
ments.' 

Fig.  202  represents 
the  method  of  manu- 
facture of  paint.  The 
body,  together  with  a 
little  oil,  enters  at  A 
and  is  ground  in  succession  in  B,  C,  D,  and  E,  during  which  proc- 
ess the  requisite  amounts  of  oil,  dryer,  and  pigment  are  added. 


FIG.  202.   Factory  appliances  used  in 
mixing  and  grinding  paint 


Storage  cell.  The  storage  cell,  or  accumulator,  plays  an 
important  part  in  modern  electrical  developments.  Its 
fundamental  characteristic  is  that  the  chemical  action  upon 
which  it  depends  is  reversible.  The  chemical  action  taking 
place  when  the  cell  is  delivering  current  is  reversed  when 
a  current  is  conducted  through  the  cell  in  an  opposite 


522    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


direction.  Electrical  energy  can  therefore  be  stored  in  the 
cell  as  chemical  energy  and  drawn  off  again,  when  desired, 
as  electrical  energy. 

In  the  ordinary  accumulator 
(Fig.  203)  the  electrodes  are 
made  of  a  skeleton  of  lead. 
When  ready  for  use  the  one 
plate  (^4)  is  covered  with  a 
thick  deposit  of  spongy  lead, 
which  is  the  active  material; 
the  other  (2?)  is  similarly 
covered  with  a  layer  of  lead 
dioxide.  The  electrolyte  is 
moderately  dilute  sulfuric  acid. 
A  number  of  pairs  of  such 
plates  are  arranged  together 
in  one  cell.  When  the  plates 
FIG.  203.  A  seven-plate  storage  are  connected  by  a  wire  the 
battery  reaction  is  as  follows:  " 


Pb  +  PbO2  +  2  H2SO4 

(Discharge >-) 


2  PbSO4  +  2  H2O 

H Charge) 


It  will  be  seen  that  the  discharge  of  the  cell  results  in 
bringing  the  two  plates  to  an  identical  condition  and  in 
withdrawing  sulfuric  acid  from  the  electrolyte.  The  cell 
is  never  allowed  to  come  entirely  to  this  discharged  con- 
dition. When  the  current  is  reversed  the  two  plates  are 
restored  to  their  original  state. 

Thorium  and  cerium.  These  elements  are  found  in  a 
few  rare  minerals,  especially  in  the  monazite  sand  of  the 
Carolinas  and  Brazil.  The  oxides  of  these  elements  are 
used  in  the  preparation  of  the  Welsbach  mantles  for  gas 
lights  (p.  312)  because  of  the  intense  light  given  out 


TIN  AND  LEAD  523 

when  a  mixture  of  the  oxides  is  heated.  These  mantles 
contain  the  oxides  of  cerium  and  thorium  in  the  ratio  of 
about  1  per  cent  of  the  former  to  99  per  cent  of  the 
latter.  Compounds  of  thorium,  like  those  of  radium,  are 
found  to  possess  radioactivity,  but  in  a  less  degree. 


1.  How  could  you  detect  lead  if  present  in  tin  foil  ? 

2.  Stannous  chloride  reduces  gold  chloride  (AuCl8)  to  gold.  Give 
the  equation. 

3.  What  are  the  products  of  hydrolysis  when  stannic  chloride  is 
used  as  a  mordant? 

4.  How  could  you  detect  arsenic  or  copper  in  lead  ? 

5.  Why  is  lead  so  extensively  used  for  making  water  pipes? 

6.  What  sulf ates  other  than  lead  sulf ate  are  insoluble  ? 

7.  Could  lead  nitrate  be  used  in  place  of  barium  chloride  in  testing 
for  sulf  ates? 

8.  How  much  lead  peroxide  could  be  obtained  from  1  kg.  of 
minium  ? 

9.  The  purity  of  white  lead  is  usually  determined  by  observing 
the  volume  of  carbon  dioxide  given  off  when  it  is  treated  with  an 
acid.   On  the  supposition  that  it  has  the  formula  2  PbCO8  •  Pb(OH)2, 
how  nearly  pure  was  a  sample  if  1  g.  gave  30  cc.  of  carbon  dioxide 
at  20°  and  750  mm.  ? 

10.  Silicon  belongs  in  the  same  family  with  tin  and  lead.    In  what 
respects  are  these  elements  similar? 

11.  What  weight  of  tin  could  be  obtained  by  the  reduction  of 
1  ton  of  cassiterite  ? 

12.  What  reaction  would  you  expect  to  take  place  when  lead 
peroxide  is  treated  with  hydrochloric  acid? 

13.  White  lead  is  often  adulterated  with  barytes.  Suggest  a  method 
for  detecting  it,  if  present,  in  a  given  example  of  white  lead. 

14.  Sugar  of  lead  may  be  prepared  by  treating  litharge  with 
acetic  acid.    What  weight  of  litharge  is  necessary  for  the  prepara- 
tion of  1000  g.  of  sugar  of  lead? 


CHAPTER  XLI 
MANGANESE  AND  CHROMIUM 


NAME 

SYMBOL 

ATOMIC 

WEIGHT 

DENSITY 

MELTING 

POINT 

FORMULAS  OF  ACIDS 

Manganese 
Chromium 

Mn 
Cr 

54.93 
52.00 

8.01 
7.3 

1260° 
1520° 

H2Mn04  and  HMnO4 
H2Cr04  and  H2Cr207 

General.  Manganese  and  chromium,  while  belonging  to 
different  families,  have  so  many  features  in  common  in 
their  chemical  conduct  that  they  may  be  studied  together 
with  advantage.  They  differ  from  most  of  the  elements 
so  far  studied  in  that  they  can  act  either  as  base-forming 
or  as  acid-forming  elements.  As  base-forming  elements  each 
of  the  metals  forms  two  series  of  salts.  In  the  one  series 
the  metal  is  bivalent,  and  in  the  other  series  it  is  trivalent. 
The  acids  in  which  these  elements  play  the  part  of  a  non- 
metal  are  unstable,  but  their  salts  are  usually  stable,  and 
some  of  them  are  important  compounds. 

Elements  like  manganese  and  chromium,  which  in  ele- 
mentary form  have  all  the  properties  of  metals  but  which 
are  capable  of  forming  acids  in  higher  states  of  valence, 
are  called  metallo-acid  elements. 

MANGANESE 

Occurrence.    Manganese  is  found  in  nature  chiefly  as  the 
dioxide  MnO2,  called  pyrolusite.     In  smaller   amounts  it 
occurs  as  the  oxides  Mn2O3  and  MngO4  and  as  the  car- 
bonate MnCOg.    Some  iron  ores  also  contain  manganese. 
524 


MANGANESE  AND  CHROMIUM  525 

Preparation  and  properties.  The  element  is  difficult  to 
prepare  in  pure  condition  and  has  no  commercial  applica- 
tions. It  can  be  prepared,  however,  by  reducing  the  oxide 
with  aluminium  powder  (p.  455)  or  by  the  use  of  the 
electric  furnace,  with  carbon  as  the  reducing  agent.  The 
metal  somewhat  resembles  iron  in  appearance,  but  is  harder, 
more  fusible,  and  more  readily  acted  upon  by  air  and  mois- 
ture. Acids  readily  dissolve  it,  forming  manganous  salts. 

By  the  reduction  of  a  mixture  of  the  oxides  of  iron  and 
manganese  an  alloy  is  obtained  known  as  ferromanganese. 
This  is  used  very  extensively  as  a  purifier  in  making  steel. 
About  14  Ib.  of  manganese  is  used  for  every  ton  of  steel 
produced. 

Oxides  of  manganese.  The  following  oxides  of  manga- 
nese are  known :  MnO,  Mn2O8,  MngO4,  MnO2,  and  Mn2O7. 
Only  one  of  these,  the  dioxide,  needs  special  mention. 

Manganese  dioxide  (pyrolusite)  (Mn02).  This  substance 
is  the  most  abundant  manganese  compound  found  in  nature, 
and  is  the  ore  from  which  all  other  compounds  of  manga- 
nese are  made.  It  is  a  hard,  brittle,  black  substance  which 
is  valuable  as  an  oxidizing  agent.  It  will  be  recalled  that 
it  is  used  in  the  laboratory  preparation  of  chlorine  and 
oxygen,  in  decolorizing  glass  which  contains  iron,  and  in 
the  manufacture  of  ferromanganese. 

Compounds  containing  manganese  as  a  base-forming  ele- 
ment. As  has  been  stated  previously  manganese  forms 
two  series  of  salts.  The  most  important  of  these  salts,  all 
of  which  belong  to  the  manganous  series,  are  the  following : 

Manganous  chloride MnCLj  •  4  H2O 

Manganous  sulfide MnS 

Manganous  sulfate MnSO4  •  4  H2O 

Manganous  carbonate MnCO3 

Manganous  hydroxide Mn(OH)a 


526    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

The  chloride  and  the  sulfate  may  be  prepared  by  heating  the 
dioxide  with  hydrochloric  and  sulf uric  acids  respectively : 

MnO2  +  4  HC1  — >•  MnCl2  +  2  H2O  +  C12 
2  MnO2  +  2  H2SO4  — >•  2  MnSO4  +  2  H2O  +  O2 

The  sulfide,  carbonate,  and  hydroxide,  being  insoluble,  may 
be  prepared  from  a  solution  of  the  chloride  or  the  sulfate 
by  precipitation  with  the  appropriate  reagents.  Most  of 
the  manganous  salts  are  rose  colored.  They  not  only  have 
formulas  similar  to  the  ferrous  salts  but  resemble  them  in 
many  of  their  chemical  properties. 

Compounds  containing  manganese  as  an  acid-forming  ele- 
ment. Manganese  forms  two  unstable  acids ;  namely, 
manganic  acid  (H2MnO4)  and  permanganic  acid  (HMnO4). 
While  these  acids  are  of  little  interest,  some  of  their  salts, 
especially  the  permanganates,  are  important  compounds. 

Manganic  acid  and  manganates.  When  manganese  diox- 
ide is  fused  with  an  alkali  and  an  oxidizing  agent  a  green 
compound  is  formed.  The  equation,  when  caustic  potash 
is  used,  is  as  follows : 

Mn02  +  2  KOH  +  [O] >-  K2MnO4  +  H2O 

The  green  compound,  K2MnO4,  is  called  potassium  man- 
ganate  and  is  a  salt  of  the  unstable  manganic  acid.  The 
manganates,  as  well  as  the  free  acid,  are  all  very  unstable. 
Permanganic  acid  and  the  permanganates.  When  potas- 
sium manganate  is  treated  with  water  a  reversible  decom- 
position takes  place  which  can  be  expressed  by  the  equation 

3  K2MnO4  +  4  H2O  <=±  2  KMnO4  +  Mn(OH)4  +  4  KOH 

If  an  acid  is  now  added  the  potassium  hydroxide  is  neu- 
tralized and  the  reaction  goes  on  to  completion,  the  chief 
product  being  potassium  permanganate,  KMnO4.  Since  one 


MANGANESE  AND  CHROMIUM  527 

third  of  the  manganese  is  lost  in  this  reaction,  potassium 
permanganate  is  now  made  by  electrolysis.  The  manganate 
is  dissolved  in  water,  and  at  the  anode  the  oxygen  evolved 
by  the  electrolysis  of  water  acts  upon  the  manganate  in 
the  following  way: 

2  K2MnO4  +  H2O  +  [O] >-  2  KMnO4  +  2  KOH 

Potassium  permanganate  crystallizes  in  dark-purple  needles 
and  is  very  soluble  in  water,  forming  an  intensely  purple 
solution.  All  other  permanganates,  as  well  as  permanganic 
acid  itself,  are  very  soluble  and  give  solutions  of  the  same 
color.  Unfortunately  sodium  permanganate  (NaMnO4)  is 
very  difficult  to  crystallize,  so  that  the  cheaper  sodium 
cannot  be  used  to  replace  potassium  in  this  salt. 

Oxidizing  properties  of  the  permanganates.  The  perman- 
ganates are  remarkable  for  their  strong  oxidizing  properties. 
When  used  as  an  oxidizing  agent  the  permanganate  is  itself 
reduced,  the  exact  character  of  the  products  formed  from  it 
depending  upon  whether  the  oxidation  takes  place  (1)  in 
an  alkaline  or  neutral  solution  or  (2)  in  an  acid  solution. 

1.  Oxidation  in  alkaline  or  neutral  solution.     When  the 
solution   is   either  alkaline   or  neutral   the  potassium  and 
the  manganese  of  the  permanganate  are  both  converted  into 
hydroxides,  as  shown  in  the  equation 

2  KMnO4  +  5  H2O  — *•  2  Mn(OH)4  +  2  KOH  +  3  [O] 

2.  Oxidation  in  acid  solution.     When  free  acid  such  as 
sulfuric  acid  is  present  the  potassium  and  the  manganese  are 
both  changed  into  salts  of  the  acid : 

2  KMnO4  +  3  H2SO4 — >•  K2SO4 +  2  MnSO4  +  3  H2O  +  5  [O] 

Under  ordinary  conditions,  however,  neither  one  of  these 
reactions  takes  place  except  in  the  presence  of  a  third 


528    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

substance  which  is  capable  of  oxidation.  The  oxygen  is 
not  given  off  in  the  free  state,  as  the  equations  show,  but 
is  used  up  in  effecting  oxidation. 

Potassium  permanganate  is  particularly  valuable  as  an 
oxidizing  agent  not  only  because  it  acts  readily  either  in 
acid  or  in  alkaline  solution  but  also  because  the  reaction 
takes  place  so  easily  that  often  it  is  not  even  necessary 
to  heat  the  solution  to  secure  action.  The  substance  finds 
many  uses  in  the  laboratory,  especially  in  analytical  work. 
It  is  also  used  as  an  antiseptic  and  a  disinfectant. 

CHROMIUM 

Occurrence.  The  ore  from  which  all  chromium  compounds 
are  made  is  chromite,  or  chrome  iron  ore  (FeCr2O4).  This 
is  found  most  abundantly  in  New  Caledonia  and  Greece, 
while  in  the  United  States  it  is  produced  in  California  and 
Oregon.  The  element  also  occurs  in  small  quantities  in 
many  other  minerals,  especially  in  crocoite  (PbCrO4),  in 
which  mineral  it  was  first  discovered. 

Preparation  and  properties.  Chromium,  like  manganese, 
is  very  hard  to  reduce  from  its  ores,  owing  to  its  great 
affinity  for  oxygen.  It  can,  however,  be  made  by  the  same 
methods  which  have  proved  successful  with  manganese. 

Chromium  is  a  very  hard  metal  of  about  the  same  den- 
sity and  melting  point  as  iron.  At  ordinary  temperatures 
air  has  little  action  on  it ;  at  higher  temperatures,  however, 
it  burns  brilliantly.  Nitric  acid  has  no  action  on  it,  but 
hydrochloric  and  dilute  sulfuric  acids  dissolve  it,  liberating 
hydrogen  and  forming  chromous  salts. 

Chromium  is  a  valuable  metal  for  producing  certain  steel 
alloys.  For  this  purpose  ferrochromium  is  first  made  by 
reducing  chromite  with  carbon,  and  this  is  added  to  the 


MANGANESE  AND  CHROMIUM  529 

liquid  steel  in  quantities  sufficient  to  produce  the  desired 
alloy.  The  metal  also  forms  desirable  alloys  with  other 
metals,  such  as  copper  and  nickel. 

Compounds  containing  chromium  as  a  base-forming  element. 
While  chromium  forms  two  series  of  salts,  chromous  salts 
are  difficult  to  prepare  in  the  solid  state  and  are  of  little 
importance.  The  most  important  of  the  chromic  series 
are  the  following: 

Chromic  hydroxide     .     !     .     .     .     .  Cr(OH)3 

Chromic  sulfate Cr2(SO4)8 

Chromic  chloride CrCl3  •  6  H2O 

Potassium  chrome  alum KCr(SO4)2  •  12  H2O 

Chromic  hydroxide,  Cr(OH)3.  This  substance  being 
insoluble  can  be  obtained  by  precipitating  a  solution  of 
the  chloride  or  sulfate  with  a  soluble  hydroxide.  It  is  a 
greenish  substance  which,  like  aluminium  hydroxide,  dis- 
solves both  in  alkalies  and  in  acids. 

Dehydration  of  chromium  hydroxide.  When  heated  gently 
chromic  hydroxide  loses  a  part  of  its  oxygen  and  hydrogen, 
forming  the  substance  CrO  •  OH,  which,  like  the  corresponding 
aluminium  compound,  has  more  pronounced  acid  properties 
than  the  hydroxide.  It  forms  a  series  of  salts  very  similar  to 
the  spinels ;  chromite  is  the  ferrous  salt  of  this  acid,  having 
the  formula  Fe(Cr02)2.  When  heated  to  a  higher  temperature 
chromic  hydroxide  is  completely  dehydrated,  forming  the  tri- 
oxide  Cr208.  This  resembles  the  corresponding  oxides  of  alu- 
minium and  iron  in  many  respects.  It  is  a  bright-green  powder, 
the  shade  of  green  depending  upon  the  exact  conditions  of 
preparation.  Both  the  hydroxide  and  the  oxide  are  used  as 
green  pigments. 

Chromic  sulfate,  Cra(S04)3.  This  compound  is  a  violet- 
colored  solid  which  dissolves  in  water,  forming  a  solu- 
tion of  the  same  color.  This  solution,  however,  turns 


530    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

green  when  heated,  owing  to  the  formation  of  basic 
salts.  Chromic  sulfate,  like  ferric  and  aluminium  sul- 
fates,  unites  with  the  sulfates  of  the  alkali  metals  to  form 
alums,  of  which  the  best  known  are  potassium  chrome 
alum,  KCr(SO4)2  •  12H2O  and  ammonium  chrome  alum, 
NH4Cr(SO4)2  •  12  H2O.  These  form  beautiful  dark-purple 
crystals  and  have  some  practical  uses  in  the  tanning 
industry  and  in  photography.  A  number  of  the  salts  of 
chromium  are  also  used  in  the  dyeing  industry,  for  they 
hydrolyze  like  aluminium  salts  and  the  hydroxide  forms 
a  good  mordant. 

Hydrolysis  of  chromium  salts.  When  ammonium  sulfide  is 
added  to  a  solution  of  a  chromium  salt,  such  as  the  sulfate, 
chromium  hydroxide  is  precipitated  instead  of  the  sulfide.  This 
is  due  to  the  fact  that  chromic  sulfide,  like  aluminium  sulfide, 
hydrolyzes  in  the  presence  of  water,  forming  chromic  hydroxide 
and  hydrogen  sulfide.  Similarly,  a  soluble  carbonate  precipi- 
tates a  basic  carbonate  of  chromium. 

Compounds  containing  chromium  as  an  acid-forming  ele- 
ment. Like  manganese,  chromium  forms  two  unstable 
acids,  namely,  chromic  acid  and  dichromic  acid.  Their 
salts,  the  chromates  and  dichromates,  are  important 
compounds. 

Chromates.  When  a  chromium  compound  is  fused  with 
an  alkali  and  an  oxidizing  agent  a  chromate  is  produced. 
When  sodium  hydroxide  is  used  as  the  alkali  the  equa- 
tion is 

2  Cr(OH)8  +  4  NaOH  +  3  [O] *•  2  Na2CrO4  +  5  H2O 

This  reaction  recalls  the  formation  of  a  manganate  under 
similar  conditions. 

Properties  of  chromates.  The  chromates  are  salts  of  the 
unstable  chromic  acid  (H2CrO4)  and  as  a  rule  are  yellow 


MANGANESE  AND  CHROMIUM  531 

in  color.  Potassium  chromate  ,(K2CrO4)  and  sodium 
chromate  (Na2GrO4)  are  freely  soluble  in  water,  but 
most  of  the  chromates  are  insoluble  and  can  be  prepared 
by  precipitation.  Thus,  when  a  solution  of  potassium 
chromate  is  added  to  a  solution  of  lead  nitrate  or  of 
barium  nitrate,  the  reactions  expressed  by  the  following 
equations  occur: 

Pb(NO3)2  +  K2CrO4  — >•  PbCrO4  +  2  KNO8 
Ba(NO8)2  +  K2Cr04 *  BaCrO4  +  2  KNO8 

Lead  chromate  (chrome  yellow)  and  barium  chromate 
separate  as  yellow  precipitates.  The  presence  of  either  of 
these  two  metals  can  be  detected  by  taking  advantage  of 
these  reactions. 

Bichromates.  When  sodium  chromate  or  potassium 
chromate  is  treated  with  an  acid,  a  salt  of  the  unstable 
dichromic  acid  (H2Cr2O7)  is  formed: 

2  Na2Cr04  +  H2SO4  — >  Na2O2O7  +  Na2SO4  +  H2O 

The  relation  between  the  chromates  and  dichromates  is 
the  same  as  that  between  the  phosphates  and  the  pyro- 
phosphates.  Sodium  dichromate  might  therefore  be  called 
sodium  pyrochromate. 

Until  recent  years  potassium  dichromate  (K2O2O7)  has 
been  the  best-known  dichromate  and  the  most  famil- 
iar chromium  compound.  It  forms  large  crystals  of  a 
brilliant-red  color  and  is  rather  sparingly  soluble  in  water. 
Sodium  dichromate  is  much  more  soluble  in  water  than  the 
potassium  salt,  and  the  crystals  deposited  from  the  satu- 
rated solution  are  not  so  well  formed  and  are  somewhat 
deliquescent.  They  have  the  formula  Na2Cr2O_  •  2  H2O. 
The  sodium  salt  is  now  more  widely  used  than  the 
potassium  salt. 


532    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Properties  and  uses  of  dichroxnates.  When  a  soluble 
dichromate  is  treated  with  a  solution  of  a  base  the 
dichromate  is  converted  into  the  chromate: 

Na2Cr2O7  +  2  NaOH  -  >-  2  Na2CrO4  +  H2O 

When  a  solution  of  a  dichromate  is  added  to  a  solution 
of  a  lead  or  a  barium  salt  the  corresponding  chromate 
(not  dichromate)  is  precipitated.  With  barium  nitrate 
the  equation  is 


—  >-  2  BaCrO4  +  2  NaNO8  +  2  HNO3 

Sodium  dichromate  finds  use  in  many  industries  as  an 
oxidizing  agent,  especially  in  the  preparation  of  organic 
substances,  such  as  the  dye  alizarin.  It  is  also  used  in 
preparing  pigments  and  in  the  process  of  tanning. 

Oxidizing  action  of  chromates  and  dichromates.  When  a 
solution  of  a  chromate  or  of  a  dichromate  is  treated  with  an 
excess  of  an  acid,  chromic  acid  is  set  free.  With  sodium 
dichromate  and  sulfuric  acid  the  equation  is  as  follows  : 

Na2Cr207  +  H2SO4  +  H2O  -  ^  Na2SO4  +  2  H2CrO4 

If  now  there  is  some  material  present  that  is  easily  oxi- 
dized, the  chromic  acid  acts  as  an  oxidizing  agent,  the 
chromium  being  converted  into  a  trivalent  metal  in  the 
process.  In  the  presence  of  sulfuric  acid  the  equation 
may  be  written  thus: 

2  H2CrO4  +  3  H2SO4  —  >•  Cr2(SO4)3  +  5  H2O  +  3  [O] 

The  oxygen  is  not  given  off  as  such,  but  is  taken  up  by 
the  substance  undergoing  oxidation.  For  example,  with 
ferrous  sulfate  the  equation  is  as  follows: 

6  FeS04  +  3  H2S04  +  3  [O]  —  ^3  Fe2(SO4)8  +  3  H2O 


MANGANESE  AND  CHROMIUM  533 

The  three  equations  showing  the  steps  in  the  oxidizing 
action  of  sodium  dichromate  upon  ferrous  sulfate  may  be 
combined  into  one  equation  : 

Na2O2O7  +  7  H2SO4  +  6  FeSO4 

—  *  Na2S04  +  Cr2(S04)3  +  3  Fe2(SO4)3  +  7  H2O 

It  will  be  seen  that  a  dichromate  decomposes  in  very  much 
the  same  way  that  a  permanganate  does,  the  sodium  and 
chromium  being  both  changed  into  salts  in  which  they  play 
the  part  of  metals,  while  part  of  the  oxygen  is  liberated. 

Potassium  chrome  alum.  When  potassium  dichromate  and 
sulfuric  acid  are  used  for  oxidation  it  will  be  noticed  that 
potassium  sulfate  and  chromium  sulfate  are  formed  as  the 
products  of  the  reaction.  On  evaporating  the  solution  these 
substances  crystallize  out  as  potassium  chrome  alum,  which 
substance  is  produced  as  a  by-product  in  the  industries  using 
potassium  dichromate  for  oxidizing  purposes.  Sodium  does  not 
form  a  similar  alum. 

Chromic  anhydride  (CrOj).  When  concentrated  sulfuric 
acid  is  added  to  a  concentrated  solution  of  potassium 
dichromate,  and  the  liquid  is  allowed  to  stand,  deep-red 
needle-shaped  crystals  appear  which  have  the  formula 
CrO3.  This  oxide  of  chromium  is  called  chromic  anhydride 
since  it  combines  readily  with  water  to  form  chromic  acid  : 

Cr08  +  H20  —  >-H2Cr04 

It  is  therefore  analogous  to  sulfur  trioxide,  which  forms 
sulfuric  acid  in  a  similar  way: 


Chromic  anhydride  is  a  very  strong  oxidizing  agent,  giving 
up  oxygen  and  forming  chromic  oxide  : 

2CrO8 


534    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


1.  How  does  pyrolusite  effect  the  decolorizing  of  glass  contain- 
ing iron? 

2.  Write  the  equations  for  the  preparation  of  manganous  chloride, 
manganous  carbonate,  and  manganous  hydroxide. 

3.  Write  the  equations  representing  the  reactions  which  take 
place  when  ferrous  sulfate  is  oxidized  to  ferric  sulfate  by  potas- 
sium permanganate  in  the  presence  of  sulfuric  acid. 

4.  In  the  presence  of  sulfuric  acid,  oxalic  acid  is  oxidized  by 
potassium  permanganate  according  to  the  equation 

C2H2O4  +  O >  2  CO2  +  H2O 

Write  the  complete  equation. 

5.  10  g.  of  iron  were  dissolved  in  sulfuric  acid  and  oxidized  to 
ferric   sulfate  by  potassium   permanganate.    What  weight  of  the 
permanganate  was  required? 

6.  What  weight  of  ferrochromium  containing  40%  chromium 
must  be  added  to  a  ton  of  steel  to  produce  an  alloy  containing  1%  of 
chromium  ? 

7.  Write  the  equation  representing  the  action  of  ammonium 
sulfide  upon  chromium  sulfate. 

8.  Sodium  chromate  oxidizes  hydrochloric  acid,  forming  chlorine. 
Write  the  complete  equation. 

9.  Give  the  action  of  sulfuric  acid  on  sodium  dichromate  («)  in 
the  presence  of  a  large  amount  of  water ;  (6)  in  the  presence  of  a 
small  amount  of  water. 

10.  What  would  be  the  percentage  composition  of  ferrochromium 
made  by  reducing  pure  chromite  ? 

11.  What  per  cent  of   the  total  oxygen  present  in  potassium 
dichromate  is  available  as  an  oxidizing  agent? 


CHAPTER  XLII 
URANIUM ;  RADIUM  AND  THORIUM 

Uranium.  Uranium  is  a  rare  element  whose  compounds 
were  first  isolated  from  a  mineral  called  pitchblende  or 
uraninite,  which  is  essentially  an  oxide  of  the  formula 
U3O8.  Carnotite,  a  mineral  discovered  more  recently  (1899), 
contains  both  uranium  and  vanadium  and  is  found  chiefly 
in  Colorado  and  Utah.  The  carnotite  ores  are  by  far  the 
most  abundant  source  of  uranium,  the  next  in  importance 
being  the  pitchblende  deposits  in  Bohemia  and  Saxony. 
The  American  production  of  uranium  ores  is  much  greater 
than  that  of  all  other  countries  combined. 

Compounds  and  uses.  The  most  familiar  compounds  of 
uranium  are  the  black  oxide,  U3Og,  uranyl  nitrate, 
UO2(NO3)2,  and  sodium  diuranate  (Na2U2O7).  Most  of 
the  compounds  of  uranium  are  yellow  or  red.  Their 
chief  chemical  use  is  in  making  greenish-yellow  fluorescent 
glass,  in  the  decorating  of  china  with  various  shades  of 
yellow,  orange,  and  black,  and  in  the  making  of  orange- 
colored  pigments.  Uranium  steel  alloy  has  useful  proper- 
ties. It  is  said  to  be  used  for  making  heavy  cannon,  and 
to  stand  the  heat  and  strain  of  rapid  fire  better  than  any 
other  steel  alloy. 

Radioactivity  of  uranium.    In  1896  the  French  physicist 

Becquerel  discovered  that  uranium  and  all  its  compounds 

possess  a  property  which  has  been  named  radioactivity. 

This  ra'dioactivity  manifests  itself  in  the  following  ways : 

635 


536    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


FIG.  204.  A  radiograph  of  some  metal  objects 


(1)  A  photographic  plate  wrapped  in  black  paper  and 
placed  near  a  compound  of  uranium  is  affected  as  though 
exposed  to  light.  A  metallic  object  placed  on  the  plate 

screens  the  plate 
from  this  action  and 
leaves  its  outline  on 
the  plate  when  it 
is  developed,  form- 
ing a  radiograph 
(Fig.  204).  (2)  A 
charged  electroscope 
is  rapidly  discharged 
when  any  material  containing  uranium  is  brought  near  it, 
which  shows  that  the  air  all  about  this  material  is  an 
electrical  conductor. 

Fig.  205  represents  a  simple  form  of  aluminium-leaf  elec- 
troscope, the  leaves  assuming  the  position  indicated  at  B  when 
an  electric  charge  is  communicated  to  the  knob  A.  When  a 
substance  containing  uranium  (C,  Fig.  206)  is  brought  near 
the  knob  A  the  charge  is  rapidly  lost,  and  the  leaves  collapse 
as  shown  at  B. 

The  discovery  of  radium.  Pitchblende 
was  found  to  be  four  times  as  radio- 
active as  uranium  itself,  which  sug- 
gested that  possibly  there  was  some 
unknown  element  in  the  mineral  that 
was  carried  over  into  the  uranium  salts 
as  an  impurity  and  was  responsible  for 
the  radioactivity.  Accordingly  Monsieur 
and  Madame  Curie  worked  over  the  residues  from  a  large 
quantity  of  pitchblende,  and  obtained  a  minute  quantity  of 
the  chloride  of  a  new  element,  which  they  named  radium. 
This  chloride  is  3,000,000  times  as  active  as  uranium. 


FIG.  205.   A  charged 
electroscope 


URANIUM;  RADIUM  AND  THORIUM         537 


FIG.  206.  A  discharged 
electroscope 


The  atomic  weight  of  radium  is  226.0,  and  this  weight, 

as  well  as  all  of  the  other  properties  of  the  element  and 
of  its  compounds,  place  it  in  the  cal- 
cium family,  just  below  barium.  The 
metal  itself  was  isolated  by  Madame 
Curie  (Fig.  207)  in  1910  and  is  very 
similar  to  barium. 

Quantity  of  radium  available.  Know- 
ing the  radioactivity  of  both  uranium 
and  radium,  it  is  not  difficult  to  esti- 
mate the  proportion  of  radium  in  any 

ore  containing  uranium.     Estimates  of  this  kind  bring  to 

light  a  very  surprising  fact 

—  the  proportion  of  radium 

in   all  classes    of   uranium 

ores  is  very  constant,  and  is 

about  1  part  of  radium  in 

2,940,000  parts  of  uranium. 

Probably    not    more    than 

50  g.    or    60  g.    of   radium 

has  been  produced.    In  the 

form  of  the  chloride,  RaCl2, 

it  costs  about  $90,000  per 

gram. 

Disintegration  of  radium. 

The  extraordinary  fact  about 

radium  is  that  although  it 

is  a  well-characterized  ele- 
ment, it  is  slowly  disinte- 
grating. In  this  process  it 

is  resolved  into  two  other 

elements,   one   of   which   is    helium  and   the   other  niton. 

These  both  belong  in  periodic  Group  0  with  the  inactive 


FIG.  207.     Madame  Curie  (1867-), 
Professor  of  physics  in  the  Univer- 
sity of  Paris 


538    AN"  ELEMENTARY  STUDY  OF  CHEMISTRY 

gases  of  the  atmosphere.  Niton,  in  turn,  decomposes  into 
helium  and  still  another  element  named  radium  A.  Similar 
decompositions  continue  through  a  number  of  stages,  and 
it  is  thought  that  the  final  product  is  lead. 

In  these  decompositions  two  distinct  kinds  of  particles 
are  shot  off  with  enormous  velocity :  (1)  the  one  kind, 
called  alpha  (a)  rays,  consists  of  helium  atoms  charged 
positively  and  moving  with  a  velocity  about  J^-  that  of 
light ;  (2)  the  other,  called  beta  (/3)  rays,  consists  of 
particles  not  more  than  jgW  of  the  weight  of  a  hydrogen 
atom  and  negatively  charged.  These  are  usually  called 
electrons,  and  their  initial  velocity  is  nearly  that  of  light. 
There  is  a  third  type  of  radiation  known  as  gamma  (7) 
rays  which  does  not  consist  of  material  particles,  but  of 
waves  in  the  ether  like  light.  The  gamma  rays  resemble 
very  closely  the  X  rays  of  the  Crookes  tubes,  but  are 
much  more  powerful. 

The  rate  at  which  the  decomposition  of  radium  proceeds 
cannot  be  changed  by  any  means  that  has  yet  been  tried. 
It  is  not  affected  by  very  high  temperature  nor  by  the 
nature  of  the  radium  compound.  * 

Demonstration  of  the  three  types  of  rays.  The  existence 
of  three  different  types  of  rays  can  be  demonstrated  in 
the  following  way:  A  small  quantity  of  material  rich 
in  radium  is  placed  at  the  bottom  of  a  hole  bored  in  a 
piece  of  lead  (^4,  Fig.  208).  A  photographic  plate  (5), 
protected  from  light,  is  fixed  at  some  distance  above  the 
radium,  and  the  poles  of  a  magnet  C,  C,  are  arranged 
one  on  each  side  of  the  hole  in  A  and  somewhat  above 
it.  The  rays  from  the  radium  are  shot  out  of  the  hole 
like  bullets  from  a  rifle.  The  positively  charged  alpha 
rays  are  bent  out  of  their  course  by  the  attraction  of 
the  negative  pole,  while  the  negatively  charged  beta  rays 


URANIUM;  RADIUM  AND  THORIUM         539 

are  bent  in  the  opposite  direction  and  to  a  much  greater 
extent  by  the  attraction  of  the  positive  pole.  The  gamma 
rays  are  not  deflected  at  all.  By  developing  the  photo- 
graphic plate  the  extent  of  deflection  of  the  alpha  and 
beta  rays  can  be  measured,  and  from  such  measurements 
it  is  possible  to  arrive  at  conclusions  as  to  the  relative 
masses  and  charges  of  the  two  kinds  of  particles. 

Origin  of  radium.     Radium  is   decomposing  at  a  rate 
which   places   its   average   life    at   2500   years,   yet  it  is 
found   in   ores   which   are  undoubtedly  much  older  than 
this.     It  must  therefore  .  . ..  ..-J-— 

be  in  the  process  of 
formation  from  some 
other  element.  Experi- 
ment leaves  no  doubt 
that  this  element  is 
uranium.  The  quantity 
of  radium  so  constantly 
present  in  ores  of  ura- 
nium  simply  represents 
the  equilibrium  between  the  rate  at  which  uranium  disinte- 
grates and  that  at  which  radium  disintegrates.  If  this  is  the 
case,  it  is  clear  that  we  can  never  hope  to  find  any  deposits 
of  radium  richer  than  those  afforded  by  uranium  ores. 

Energy  of  radium.  During  the  decomposition  of  radium, 
niton,  and  the  succeeding  products  a  very  great  deal  of 
energy  is  given  off.  Both  the  helium  atoms  and  the  elec- 
trons are  shot  off  with  very  high  kinetic  energy,  and  the 
radium  compound  is  kept  heated  by  the  heat  energy  set 
free.  It  is  estimated  that  1  gram  of  radium  hourly  evolves 
132  calories  of  heat.  From  this  value,  together  with  the 
average  life  period  (2500  years),  it  is  easy  to  compute 
that  the  total  energy  given  off  by  a  gram  of  radium  will 


540    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


be  250,000  times  the  heat  of  combustion  of  a  gram  of 
carbon.  These  unquestioned  facts  have  thrown  a  great 
deal  of  doubt  upon  the  older  estimates  of  the  age  of 
the  earth. 

Radium  and  the  atomic  conception.  It  is  clear  that  the 
atom  of  radium,  as  well  as  that  of  uranium,  must  have  a 
very  elaborate  structure,  since  helium,  electrons,  and  free 

energy  are  formed  from 
them.  All  the  facts 
in  the  case  are  brought 
into  a  fair  degree  of 
harmony  if  we  assume 
that  these  atoms  are 
made  up  (1)  of  a  posi- 
tive nucleus  consisting 
of  charged  atoms  of 
helium  (and  probably 
of  hydrogen)  closely 
packed  together  and 
constituting  practically 
all  the  mass  of  the 
atom;  and  (2)  a  number 
of  negative  electrons 
revolving  about  the  nu- 
cleus much  as  planets  revolve  about  the  central  sun. 

In  the  case  of  uranium  and  radium  the  system  is  un- 
stable for  some  reason,  and  from  time  to  time  atoms 
explode  expelling  either  electrons  or  atoms  of  helium, 
or  both,  and  forming  other  atomic  systems.  The  violent 
explosion  causes  an  intense  wave-motion  in  the  ether  that 
constitutes  the  gamma  ray.  In  the  case  of  most  of  the 
elements  the  atoms  are  stable,  but  they  are  probably 
made  up  of  systems  similar  to  the  ones  described. 


FIG.  209.  Total  quantity  of  radium  bro- 
mide (1.764g.)  extracted  from  300  tons 
of  carnotite  (actual  size) 


URANIUM;  RADIUM  AND  THORIUM         541 

This  conception,  which  has  many  facts  to  support  it, 
affords  us  some  insight  into  the  meaning  of  the  periodic 
law,  for  the  mass  of  any  given  atom  will  depend  upon 
the  number  and  character  of  the  simple  atoms  constituting 
the  nucleus. 

The  use  of  radium  in  tlie  treatment  of  disease.  The 
rays  emitted  from  radium,  niton,  and  other  radioactive 
elements  produce  many  chemical  and  physiological  effects. 
They  disintegrate  glass,  water,  and  many  other  substances. 
They  produce  severe  burns  upon  the  skin,  like  those  of 
X  rays.  They  kill  bacteria  and  other  microorganisms. 

This  latter  property  has  led  to  the  hope  that  exposure 
to  the  radiations'  of  radium  compounds  might  prove  to 
be  of  assistance  in  effecting  a  cure  for  some  diseases  of 
the  skin  and  for  cancer.  It  is  not  possible  as  yet  to  say 
to  what  extent  these  hopes  will  be  realized.  Certain 
forms  of  cancer  have  apparently  been  cured  in  this  way. 

Radioactive  thorium.  The  rare  element  thorium  exhibits 
properties  very  similar  to  those  of  uranium.  It  gives  rise 
to  the  same  kind  of  series  of  radioactive  elements  by 
successive  decomposition,  producing  the  same  varieties  of 
radiation  as  the  other  series.  Uranium  and  thorium  are 
the  elements  of  greatest  atomic  weight,  and  no  other 
common  elements  are  known  to  possess  similar  proper- 
ties. This  suggests  the  idea  that  possibly  elements  of 
still  higher  atomic  weight  may  have  existed  at  some 
time,  but  that  they  have  disintegrated  to  form  elements 
of  smaller  atomic  weight  which  are  not  radioactive. 

Disintegration  series  for  uranium  and  thorium.  Our 
knowledge  of  the  stages  in  the  disintegration  of  uranium 
and  of  thorium  is  more  extensive  than  has  been  described 
in  the  account  just  given.  .  The  diagram  (Fig.  210)  pre- 
pared by  Soddy  gives  a  more  complete  summary  of  the 


542    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Uranium 

(8  X  109  years) 


Thorium 
(4X10i»yearsO 

Mesothorium  1 
(7.9  years) 

Mesothorium  2 
(8.9  hours) 

Radiothorium 
(2.91years(?)) 

Thorium  X 
(5.35  days) 

Emanation 
(76  seconds) 

Thorium  A 
(0.203  second) 

Thorium  B 
(15.3  hours) 

Thorium  Cx 
(79  minutes) 

Thorium  C2 


Thorium  D 
(4.5  minutes) 


Thorium  E 
(Unknown) 


Radium  G 
(probably  lead) 

• 

FIG.  210.   Disintegration  series  for  uranium  and  for  thorium 

name  and  atomic  weight  of  the  products  formed,  the 
average  life  of  these  products,  and  the  character  of 
the  rays  given  out  in  the  process  of  formation. 


URANIUM;  RADIUM  AND  THORIUM         543 


1.  When  was  uranium  discovered,  and  how  did  it  get  its  name? 
(See  encyclopedia.) 

2.  For  whom  was  carnotite  named?    (See  dictionary.) 

3.  How  is  an  electroscope  charged  ?   (See  physics.) 

4.  What  is  the  meaning  of  alpha,  beta,  and  gammal 

5.  What  is  the  velocity  of  light?    (See  physics.) 

6.  How  did  thorium  get  its  name  ?  For  what  is  the  element  used  ? 

7.  What  weight  of  pitchblende  is  necessary  for  the  preparation 
of  1  g.  of  radium  ? 


CHAPTER  XLIII 
THE  PLATINUM  METALS  AND  GOLD 


NAME 

SYMBOL 

ATOMIC 

WEIGHT 

DENSITY 

HIGHEST 

OXIDE 

HIGHEST 
CHLORIDE 

MELTING 

POINT 

Ruthenium  .     . 

Ru 

101.7 

12.26 

Ru04 

RuCl4 

2450° 

Rhodium     .     . 

Rh 

102.9 

12.1 

Rh02 

RhCl3 

1950° 

Palladium   .     . 

Pd 

106.7 

11.8 

Pd02 

PdCl4 

1549° 

Iridium  .     .     . 

Ir 

193.1 

22.42 

Ir02 

IrCl4 

2350° 

Osmium  .     .     . 

Os 

190.9 

22.47 

Os04 

OsCl4 

2700° 

Platinum     .     . 

Pt 

195.2 

21.50 

'Pt02 

PtCl4 

1755° 

Gold  .... 

Au 

197.2 

19.30 

Au203 

AUC13 

1063° 

The  family.  Following  iron,  cobalt,  and  nickel  in  the 
eighth  column  of  the  periodic  table  are  two  groups  of 
three  elements  each.  The  metals  of  the  first  of  these 
groups  —  ruthenium,  rhodium,  and  palladium  —  have 
atomic  weights  near  100  and  densities  near  12.  The 
metals  of  the  other  group  —  iridium,  osmium,  and  plati- 
num—  have  atomic  weights  near  200  and  densities  near 
21.  These  six  rare  elements  have  very  similar  physical 
properties  and  resemble  each  other  chemically  not  only 
in  the  type  of  compounds  which  they  form  but  also  in 
the  great  variety  of  them.  They  occur  closely  associated 
in  nature,  usually  as  alloys  of  platinum  in  the  form  of 
irregular  metallic  grains  in  sand  and  gravel.  They  are 
known  collectively  as  the  platinum  metals.  Platinum 
and  palladium  are  by  far  the  most  abundant  of  the  six. 

Although  the  periodic  classification  assigns  gold  to  the 
silver-copper  group,  it  much  more  closely  resembles  the 
544 


THE  PLATINUM  METALS  AND  GOLD        546 

platinum  metals  in  its  physical  properties  as  well  as  in 
its  chemical  conduct,  and  it  can  be  conveniently  consid- 
ered along  with  them.  The  four  elements  gold,  platinum, 
osmium,  and  iridium  are  the  heaviest  substances  known, 
being  about  twice  as  heavy  as  lead. 

PLATINUM 

Occurrence.  About  90  per  cent  of  the  platinum  of  com- 
merce comes  from  Russia,  small  amounts  being  produced 
in  Colombia,  New  South  Wales,  Canada,  and  the  United 
States.  Owing  to  the  rapidly  increasing  demands  for  the 
metal  great  efforts  are  being  made  to  discover  new  de- 
posits and  a  number  of  less  important  ones  have  been 
announced.  Some  ores  of  gold,  copper,  and  nickel  con- 
tain very  small  percentages  of  platinum,  and  this  plati- 
num is  now  recovered  in  the  process  of  refining.  In  the 
United  States  about  24,500  oz.  is.  now  produced  annually 
from  this  source,  chiefly  from  blister  copper  and  gold 
bullion. 

Preparation.  Native  platinum  is  usually  alloyed  with 
gold  and  the  platinum  metals.  To  separate  the  platinum 
the  alloy  is  dissolved  in  aqua  regia,  which  converts  the 
platinum  into  chloroplatinic  acid  (H2PtCl6).  Ammonium 
chloride  is  then  added,  which  precipitates  the  platinum  as 
insoluble  ammonium  chloroplatinate : 

H2PtCl6  +  2  NH4C1  — >•  (NH4)2PtCl6  +  2  HC1 

Some  indium  is  also  precipitated  as  a  similar  compound. 
On  ignition  the  double  chloride  is  decomposed,  leaving 
the  platinum  as  a  spongy  metallic  mass,  which  is  melted 
in  an  electric  furnace  and  rolled  or  hammered  into  the 
desired  shape. 


546    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

Properties.  Platinum  is  a  grayish-white  metal  of  high 
luster,  and  is  very  malleable  and  ductile.  It  melts  in  the 
oxyhydrogen  blowpipe  and  in  the  electric  furnace;  it  is 
harder  than  gold  and  is  a  good  conductor  of  electricity. 
In  finely  divided  form  it  has  the  ability  to  absorb  or 
occlude  gases,  especially  oxygen  and  hydrogen.  These 
gases,  when  occluded,  are  in  a  very  active  condition  re- 
sembling the  nascent  state,  and  can  combine  with  each 
other  at  ordinary  temperatures.  A  jet  of  hydrogen  or 
coal  gas  directed  upon  spongy  platinum  is  at  once  ignited. 

Platinum  as  a  catalytic  agent.  Platinum  is  remarkable  for  its 
property  of  acting  as  a  catalytic  agent  in  a  large  number  of 
chemical  reactions,  and  mention  has  been  made  of  this  use 
of  the  metal  in  connection  with  the  manufacture  of  sulfuric 
acid.  When  desired  for  this  purpose  some  porous  or  fibrous 
substance,  such  as  asbestos,  is  soaked  in  a  solution  of  chloro- 
platinic  acid  and  then  ignited.  The  platinum  compound  is 
decomposed  and  the  platinum  deposited  in  very  finely  divided 
form.  Asbestos  prepared  in  this  way  is  called  platinized  asbes- 
tos. The  catalytic  action  seems  to  be  in  part  connected  with 
the  property  of  absorbing  gases  and  rendering  them  nascent. 
Some  other  metals  possess  this  same  power,  notably  palladium, 
which  is  remarkable  for  its  ability  to  absorb  hydrogen. 

Chemical  properties.  Platinum  is  a  very  inactive  ele- 
ment chemically,  and  is  not  attacked  by  any  of  the 
common  acids.  Aqua  regia  slowly  dissolves  it,  forming 
platinic  chloride  (PtCl4),  which  in  turn  unites  with  the 
hydrochloric  acid  present  in  the  aqua  regia,  forming  the 
compound  chloroplatinic  acid  (H2PtCl6).  Platinum  is 
slowly  attacked  by  fused  alkalies.  It  combines  at  higher 
temperatures  with  carbon  and  phosphorus  and  alloys 
with  many  metals.  It  is  readily  attacked  by  chlorine, 
but  not  by  oxidizing  agents  unless  an  alkali  is  present. 


THE  PLATINUM  METALS  AND  GOLD        647 

Applications.  The  applications  of  platinum  in  the 
sciences  and  the  industries  depend  largely  upon  its  high 
melting  point,  its  chemical  inactivity,  and  its  malleability 
and  ductility.  It  is  extensively  used  (1)  in  scientific 
laboratories  and  in  certain  industries  for  evaporating 
pans,  catalytic  materials,  and  a  great  variety  of  labora- 
tory appliances;  (2)  in  electrical  apparatus  for  contact 
points  and,  until  recently,  for  lead-wires  in  the  base  of 
incandescent  lamps ;  (3)  in  dentistry  as  pins  for  artificial 


FIG.  211.    Some  laboratory  utensils  made  of  platinum 

teeth  and  as  foil  and  plates  in  construction  work ;  (4)  in 
jewelry  as  a  substitute  for  gold.  Unfortunately  the  supply 
of  the  metal  is  very  limited,  and  its  cost  is  steadily  ad- 
vancing, so  that  it  is  now  much  more  valuable  than  gold. 
Platinum  substitutes.  To  meet  the  serious  shortage  of 
platinum,  efforts  are  being  made  to  discourage  its  use  in 
jewelry  and  to  provide  substitutes  for  other  uses.  Tung 
sten  is  taking  its  place  for  many  purposes,  especially  for 
electrical  spark  contacts ;  wires  of  nickel  iron  coated  with 
copper  are  now  used  in  the  manufacture  of  incandescent 
lamps ;  electrical-resistance  heaters  are  made  of  alloys  of 
nickel  and  chromium ;  quartz  dishes  are  substituted  for 
platinum  in  the  industries:  and  a  variety  of  alloys  are 


548    AN  ELEMENTARY  STUDY  OF  CHEMISTEY 

being  proposed  for  use  in  small  laboratory  utensils.  Among 
these  the  following  are  the  most  promising  at  present: 

Palau  :  an  alloy  of  palladium  (Pal)  and  gold  (Au) 

Rhotanium :  an  alloy  of  gold,  palladium,  and  a  small  percentage  of 

rhodium 
Stellite :  essentially  an  alloy  of  cobalt  and  chromium  with  some 

molybdenum  or  tungsten 
Illium :     essentially    an    alloy    of    nickel,    chromium,   copper,    and 

molybdenum 

The  need  of  these  substitutes  can  be  seen  from  the 
following  table  of  costs,  which  is,  of  course,  subject  to 
much  change: 

Platinum,  per  gram $3.72 

Palau,  per  gram 2.00 

Gold,  per  gram 0.70 

Compounds.  Platinum  forms  two  series  of  salts  of  which 
platinous  chloride  (PtCl2)  and  platinic  chloride  (PtCl4) 
are  examples.  Platinates  are  also  known.  While  a  great 
variety  of  compounds  of  platinum  have  been  made,  the 
substance  is  chiefly  employed  in  the  metallic  state. 

Chloroplatinic  acid  (H2PtCl6).  When  platinum  is  dis- 
solved hi  aqua  regia  and  the  solution  is  crystallized, 
brownish-red  crystals  of  the  hydrate  of  chloroplatinic 
acid  are  obtained  of  the  composition  H2PtCl6  •  6  H2O.  The 
potassium  and  ammonium  salts  of  this  acid  are  nearly 
insoluble  in  water  and  alcohol.  The  acid  is  therefore  used 
as  a  reagent  to  precipitate  potassium  in  analytical  work. 
With  potassium  chloride  the  equation  is 

2  KC1  +  H2PtCl6 >-  K2PtCl6  +  2  HC1 

Other  metals  of  the  family.  Of  the  other  members  of 
the  family  palladium  is  the  most  abundant.  It  occurs  in 
considerable  quantities  alloyed  with  platinum  in  Brazil.  It 


THE  PLATINUM  METALS  AND  GOLD        549 

is  also  present  in  minute  quantities  in  many  ores  of  copper 
and  nickel,  and  at  present  the  chief  source  of  palladium  is 
the  electrolytic  muds  of  the  nickel  refineries. 

Palladium  is  only  about  half  as  heavy  as  platinum,  melts 
much  lower,  and  is  harder.  It  is  used  as  a  solder  for 
platinum,  for  making  graduated  scales  in  scientific  instru- 
ments, for  making  alloys,  and  as  a  substitute  for  platinum 
in  jewelry.  In  the  form  of  a  powder  it  is  a  remarkably 
active  catalytic  agent. 

Iridium  gives  a  very  hard  alloy  with  platinum,  used 
for  pen  points,  compass  bearings,  and  standard  weights 
and  measures. 

Osmium  tetroxide  (OsO4)  is  a  very  volatile  liquid  and 
is  used  under  the  name  of  osmic  acid  as  a  stain  for  sections 
in  microscopy. 

GOLD 

Occurrence.  Gold  has  been  found  in  many  localities,  the 
most  famous  being  South  Africa,  Australia,  Russia,  and 
the  United  States.  In  this  country  it  is  found  in  Alaska 
and  in  nearly  half  the  states  of  the  Union  —  notably  in 
California,  Colorado,  and  Nevada.  It  is  usually  found 
in  the  native  condition,  frequently  alloyed  with  silver;  in 
combination  it  is  sometimes  found  as  telluride  (AuTe2). 
The  United  States  produces  over  one  fifth  of  the  world's 
annual  output. 

Mining  and  extraction.  Native  gold  occurs  in  the  form 
of  small  grains  or  larger  nuggets  in  the  sands  of  old 
rivers  or  imbedded  in  quartz  veins  in  rocks.  In  the  first 
case  it  is  obtained  in  crude  form  by  placer  mining.  The 
sand  containing  the  gold  is  shaken  or  stirred  in  troughs 
of  running  waters  called  sluices.  This  sweeps  away  the 
sand  but  allows  the  heavier  gold  to  sink  to  the  bottom 


550    AN  ELEMENTAEY  STUDY  OF  CHEMISTRY 

of  the  sluice.  Sometimes  the  sand  containing  the  gold  is 
washed  away  from  its  natural  location  into  the  sluices  by 
powerful  streams  of  water  delivered  under  pressure  from 
pipes.  This  is  called  hydraulic  mining.  In  vein  mining 
the  gold-bearing  quartz  is  mined  from  the  veins,  stamped 
into  fine  powder  in  stamping  mills,  and  the  gold  is  extracted 
by  a  number  of  processes,  two  of  which  will  be  described. 

1.  Amalgamation  process.    In  the  amalgamation  process 
the  powder  containing  the  gold  is  washed  over  a  series 
of  copper  plates  whose  surfaces  have  been  amalgamated 
with  mercury.    The  gold  sticks  to  the  mercury  or  alloys 
with  it,  and  after  a  time  the  gold  and  mercury  are  scraped 
off  and  the  mixture  is  distilled.    The  mercury  distills  off 
and  the  gold  is  left  in  the  retort  ready  for  refining. 

2.  Cyanide  process.    This  process  depends  upon  the  fact 
that  gold  is  soluble  in  a  solution  of  sodium  cyanide  in 
the  presence  of  the  oxygen  of  the  air.    The  powder  from 
the  stamping  mills  is  treated  with  a  very  dilute  sodium 
cyanide  solution  which  extracts  the  gold : 

4Au  +  8NaCN+2H20  +  02 MNaOH  +  4NaAu(CN)s 

From  this  solution  the  gold  can  be  obtained  by  electrolysis 
or  by  precipitation  with  metallic  zinc : 

2  NaAu(CN)2  +  Zn *•  Na2Zn(CN)4  +  2  Au 

Refining   of   gold.      Gold    is    refined   by   three    general 
methods : 

1.  Electrolysis.    When  gold  is  dissolved  in  a  solution  of 
potassium  cyanide,  and  the  solution  electrolyzed,  the  gold 
is  deposited  in  very  pure  condition  on  the  cathode. 

2.  Cupellation.     When  the  gold  is   alloyed  with  easily 
oxidizable  metals,  such  as  copper  or  lead,  it  may  be  refined 
by  cupellation.    The  alloy  is  fused  with  an  oxidizing  flame 


THE  PLATINUM  METALS  AND  GOLD        551 

on  a  shallow  hearth  made  of  bone  ash,  which  substance 
has  the  property  of  absorbing  metallic  oxides  but  not  the 
gold.  Any  silver  which  may  be  present  remains  alloyed 
with  the  gold. 

3.  Parting  with  sulfuric  acid.  Gold  may  be  separated  from 
silver,  as  well  as  from  many  other  metals,  by  heating  the 
alloy  with  concentrated  sulfuric  acid.  This  dissolves  the 
silver,  while  the  gold  is  not  attacked.  Sometimes  nitric 
acid  is  used  instead  of  sulfuric  acid. 

Physical  properties.  Gold  is  a  very  heavy  bright-yellow 
metal,  exceedingly  malleable  and  ductile,  and  a  good  con- 
ductor of  electricity.  It  is  quite  soft  and  is  usually  alloyed 
with  copper  or  silver  to  give  it  the  hardness  required  for 
most  practical  uses.  The  degree  of  fineness  is  expressed 
in  terms  of  carats,  pure  gold  being  twenty-four  carats ; 
the  gold  used  for  jewelry  is  usually  eighteen  carats, 
eighteen  parts  being  gold  and  six  parts  copper  or  silver. 
Gold  coinage  is  90  per  cent  gold  and  10  per  cent  copper. 

Chemical  conduct.  Gold  is  not  attacked  by  any  of 
the  common  acids ;  aqua  regia  easily  dissolves  it,  forming 
chlorauric  acid  (HAuCl4).  Fused  alkalies  also  attack  it. 
Most  oxidizing  agents  are  without  action  upon  it,  and  in 
general  it  is  not  an  active  element. 

Compounds.  The  compounds  of  gold,  though  numerous  and 
varied  in  character,  are  of  comparatively  little  importance  and 
need  not  be  described  in  detail.  The  element  forms  two  series 
of  salts  in  which  it  acts  as  a  metal :  in  the  aurous  series  the 
gold  is  univalent,  the  chloride  having  the  formula  AuCl ;  in 
the  auric  series  it  is  trivalent,  auric  chloride  having  the  formula 
AuCl8.  Gold  also  acts  as  an  acid-forming  element,  forming  such 
compounds  as  potassium  aurate  (KAu02).  Its  compounds  are 
very  easily  decomposed,  however,  metallic  gold  separating 
from  them. 


552    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


1.  From  the  method  of  preparation  of  platinum,  what  metal  is 
likely  to  be  alloyed  with  it  ? 

2.  The  "  platinum  chloride  "  of  the  laboratory  is  made  by  dissolv- 
ing platinum  in  aqua  regia.    What  is  the  compound? 

'  3.  How  should  you  expect  potassium  aurate  and  platinate  to  be 
formed  ?  What  precautions  would  this  suggest  in  the  use  of  platinum 
vessels  ? 

4.  Why  cannot  gold  be  used  for  laboratory  vessels  instead  of 
platinum  ? 

5.  What  uses  of  platinum  would  seem  to  permanently  lose  it? 

6.  If  platinum  and  palladium  cost  the  same  per  ounce,  which 
would  be  the  more  economical  to  use  for  jewelry? 

7.  An  18-carat  gold  ring  weighs  5  g.   What  weight  of  chlorauric 
acid  can  be  made  from  this  ? 


CHAPTER  XLIV 
SOME  APPLICATIONS  OF  THE  RARER  ELEMENTS 

Rarer  elements.  A  large  number  of  elements  are  known 
which  have  not  been  described  in  the  foregoing  pages  be- 
cause an  acquaintance  with  them  is  not  at  all  necessary 
for  an  understanding  of  the  principles  of  chemistry. 

Some  of  these,  while  comparatively  rare,  could  be  pro- 
duced in  considerable  quantities  if  there  were  any  commer- 
cial use  for  them.  A  good  example  is  tellurium,  an  element 
in  the  sulfur  family  obtained  as  a  by-product  in  copper 
refining.  Others  of  these  elements  are  so  rare  that  the  cost 
of  production  is  prohibitive,  even  though  they  have  very 
useful  properties. 

Application  in  the  industries.  Some  of  these  less  familiar 
elements  or  their  compounds  have  properties  which  make 
them  valuable  for  special  purposes,  and  mention  of  a  few 
of  these  applications  will  be  of  interest. 

The  rare  earths  constitute  a  group  of  about  sixteen 
elements,  all  trivalent  and  resembling  aluminium  in  a 
general  way.  They  are  very  difficult  to  separate  from 
each  other  and  always  occur  together  in  nature.  Very 
large  quantities  of  a  mixture  of  them  accumulate  in  the 
extraction  of  thorium  from  monazite  sand  (p.  312).  The 
only  one  whose  compounds  are  obtained  pure  rather  easily 
is  cerium.  Compounds  of  cerium  are  used  as  mordants,  as 
catalytic  agents,  and  in  medicine  and  photography.  An 
alloy  of  cerium  with  iron  is  used  as  a  gas  or  cigar  lighter, 
563 


554    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 

since  it  gives  off  a  stream  of  sparks  when  scratched  by 
hard  iron.  Mixed  salts  of  the  other  rare  earths  have  been 
used  as  antiseptics,  as  mordants,  and  as  pigments. 

Thorium  oxide,  mixed  with  1  per  cent  of  cerium  oxide, 
constitutes  the  material  of  which  most  gas  mantles  are 
made  (p.  312). 

Titanium  in  the  silicon  family  is  not  a  very  rare  element, 
occurring  chiefly  as  the  oxide  TiO2,  called  rutile,  and  as  a 
constituent  of  certain  iron  ores  (ilmenite^.  Large  quanti- 
ties of  nearly  pure  titanium  or  of  ferrotitanium  are  used 
in  making  steel  rails  designed  to  stand  very  heavy  wear 
(railway  curves  and  terminals).  Titanium  oxide  is  also 
incorporated  in  electric-arc  carbons  (flaming  arc).  Carbons 
thus  made  give  a  more  diffused  and  efficient  light  than 
those  made  from  pure  carbon.  The  oxide  is  also  used  to 
impart  a  yellow  color  to  glazes,  to  porcelain,  and  to 
artificial  teeth. 

Zirconium  oxide  is  finding  extensive  use  in  the  manufac- 
ture of  fire-brick  linings  for  furnaces  and  of  tubes  in  which 
substances  may  be  heated  to  a  high  temperature.  Zircite 
contains  86  per  cent  ZrO2,  and  zirconalba  99  per  cent. 

Vanadium  occurs  in  considerable  quantities  in  carnotite 
(p.  535)  and  in  certain  sulfides  found  in  Peru.  It  is  found 
as  the  oxide  in  the  ash  of  nearly  all  anthracite  coal. 
Ferrovanadium,  like  ferrotitanium,  is  used  in  producing 
special  grades  of  steel,  particularly  when  great  toughness 
is  desired  (automobile  parts).  Its  compounds  are  used  as 
photographic  developers,  as  catalytic  reagents  in  the  dye 
industry  (aniline  black),  as  coloring  materials  in  glass, 
and  as  mordants. 

Molybdenum  compounds  are  used  in  coloring  pottery 
and  in  dyeing  silk,  wool,  and  leather.  Ammonium  molyb- 
date  is  an  important  reagent  in  the  analysis  of  phosphates. 


APPLICATIONS  OF  THE  RARER  ELEMENTS   555 


Ferromolybdenum  has  been  used  in  making  steel  alloys, 
but  does  not  seem  to  be  superior  to  the  cheaper 
ferrotungsten. 

Tungsten  compounds  are  produced  in  fairly  large  quan- 
tities. It  has  been  found  possible  to  draw  the  metal  into 
very  fine  wire  (0.3  mm.),  which  is  now  extensively  used 
instead  of  carbon  as  a  filament  for  incandescent  lamps 
(Fig.  212).  Its  melting  point  is  very 
high  (3000°),  and  the  consumption  of 
electrical  energy  for  a  given  candle  power 
is  so  low  that  the  lamp  is  about  three 
times  as  efficient  as  the  older  (carbon) 
lamp.  The  metal  is  rapidly  replacing 
platinum  for  electrical  contacts  in  switches, 
telephone  jacks,  and  automobile  vibrators. 
Ferrotungsten  is  used  in  making  steel 
designed  for  lathe  tools,  since  such  steel 
can  be  heated  to  a  red  glow  without 
losing  temper. 

Compounds  of  tungsten  are  used  for  making  fireproof 
cloth,  and  pigments  for  paints  and  pottery,  and  as  mordants. 

Selenium,  an  element  in  the  sulfur  family,  is  obtained 
as  a  by-product  in  refining  copper.  It  is  a  nonconductor 
of  electricity  when  in  the  dark,  but  becomes  a  fairly  good 
conductor  when  exposed  to  light.  This  has  led  to  its  use  in 
automatic  fire  alarms  and  for  regulating  automatic  gas 
buoys  at  sea.  Added  to  glass,  it  produces  a  fine  red  color, 
such  glass  being  used  for  railway  lanterns.  It  is  also  used 
to  produce  red  enamels. 


FIG.  212.   A  tung- 
sten lamp 


APPENDIX 


CRYSTALLOGRAPHY 

Crystals.  When  a  liquid  freezes  it  changes  into  a  mass  of 
solid  bodies,  each  of  which  has  a  definite  geometric  form  and 
is  known  as  a  crystal.  Similar  bodies  may  also  be  deposited 
from  solutions  or  be  formed  by  condensing  vapors.  Crystals 
are  always  bounded  by  plane  surfaces,  which  are  arranged  in 
an  orderly  fashion  with  reference  to  imaginary  lines  drawn 
through  the  crystal  and  called  its  axes.  Every  crystal  has 
therefore  a  definite  geometric  form.  While  the  variety  of  form 
which  crystals  may  assume  is  almost  endless,  it  has  been  found 
that  they  may  all  be  referred  to  one  of  six  fundamental  arrange- 
ments of  axes,  these  constituting  what  are  known  as  the  systems 
of  crystallography.  These  arrangements,  together  with  two  of 
the  simplest  crystal  forms  referred  to  each,  are  shown  in  the 
accompanying  figures  (Figs.  213-219). 


FIG. 213' 


FIG.  214 


The  crystal  systems.    The  relation  of  the  axes  in  the   several 
systems  is  as  follows  : 

1.  Isometric,  or  regular,  system  (Fig.  213)  :  three  equal  axes  all  at 
right  angles  to  each  other. 

2.  Tetragonal  system  (Fig.  214) :  two  equal  axes  and  a  third  of 
different  length,  all  at  right  angles. 

557 


558    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


3.  Orthorhombic  system  (Fig.  215)  :   three  unequal  axes  all  at 
right  angles. 

4.  Monoclinic  system  (Fig.  216)  :  two  axes  at  right  angles  and  a 
third  at  right  angles  to  one  of  these  but  inclined  toward  the  other. 
The  axes  may  be  of  any  relative  lengths,  and  the  angle  of  inclination 
may  vary  from  0°  to  90° ;  but  in  a  given  substance  both  are  constant. 


FIG. 215 


FIG. 216 


5.  Triclinic  system  (Fig.  217)  :  three  axes,  all  inclined  toward 
each  other.    The  axes  may  be  of  any  relative  length,  and  the  angles 
of  inclination  may  also  vary. 

6.  Hexagonal  system  (Fig.  218)  :  three  equal  axes  in  the  same 
plane,  intersecting  at  angles  of  60°,  and  a  fourth  at  right  angles  to 
all  of  these.    In  addition  to  the  two  general  forms  shown  in  Fig.  218 
there  are  many  rhombic  forms  belonging  to  this  system  (Fig.  219), 
and  these  are  sometimes  considered  to  constitute  a  seventh  system. 


FIG.  217 


FIG.  218 


FIG. 219 


Structure  of  crystals.  There  is  little  doubt  that  these  plans  of 
formation  correspond  to  orderly  arrangements  of  the  particles  of 
solid  matter  of  which  the  crystals  are  composed,  so  that  the  crystals 
resemble  in  structure  the  piles  of  cannon  balls  in  a  military  park. 
In  accordance  with  this  idea  it  is  known  that  the  various  properties 
of  the  crystal,  such  as  hardness,  strength,  optical  refraction,  and 


APPENDIX  55d 

conductivity  toward  heat  and  electricity,  differ  in  different  directions 
through  the  crystal.  Crystals  also  split  in  definite  directions,  giving 
plane  surfaces.  Some  of  these  groupings  represent  a  more  stable 
arrangement  than  do  others,  so  that  when  a  given  substance  crystal- 
lizes in  two  forms,  as  sometimes  happens,  the  change  from  the  one 
to  the  other  is  in  general  accompanied  by  an  energy  change.  It  is 
evident  that  a  body  like  glass  might  be  cut  and  polished  so  as  to  be 
an  exact  copy  of  a  crystal,  but  would  really  not  be  one  at  all,  since 
it  would  have  none  of  the  structure  of  a  crystal. 

Crystal  form  a  characteristic  of  a  substance.  In  general,  under 
the  same  conditions,  a  given  substance  will  always  crystallize 
in  a  form  which  may  be  referred  to  the  same  system  and  with 
the  same  ratio  of  axis  lengths,  and  degree  of  inclination.  The 
actual  crystal  form  may  be  quite  different,  howerer.  For 
example,  the  form  may  be  either  a  cube  or  an  octohedron,  both 
of  which  are  referred  to  the  same  axes.  Not  infrequently  a 
substance  may,  under  different  conditions,  assume  two  forms 
in  entirely  different  systems,  and  it  is  then  said  to  be  dimor- 
phous. For  example,  one  form  may  occur  when  the  substance 
freezes,  and  another  when  it  is  deposited  from  solution. 
Trimorphous  substances  are  also  known.  When  two  substances 
crystallize  in  the  same  form  and  have  the  same  inclination  of 
axes  and  the  same  ratios  in  their  lengths,  they  are  said  to  be 
isomorphous. 

Growth  of  crystals.  When  a  substance  starts  to  crystallize  it 
usually  happens  that  many  small  crystals  form  almost  simul- 
taneously. These  small  crystals  then  increase  in  size,  and  by 
choosing  favorable  conditions  very  large  crystals  may  be 
secured.  Slow  growth  favors  the  formation  of  large  crystals, 
while  rapid  crystallization  produces  many  small  ones.  The 
crystal  is  at  all  times  in  equilibrium  with  the  liquid,  some 
molecules  dissolving  from  the  crystal  and  others  depositing 
upon  it.  For  this  reason  an  imperfect  crystal  often  becomes 
perfect  even  though  its  weight  is  unchanged. 


560    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


TABLES 

TENSION  OF  AQUEOUS  VAPOR  AT  VARIOUS  TEMPERATURES, 
EXPRESSED  IN  MILLIMETERS  OF  MERCURY 


TEMPERATURE 
0°  .... 
16°  .... 
17°  . 


PRESSURE       TEMPERATURE 


19° 


.     4.6 
.  13.62 
.  14.4 
.  15.46 
.  16.45 


21° 18.62 

22° 19.79 

23° 21.02 

24° 22.32 

25°  .  23.69 


20° 17.51       100° 


760.00 


WEIGHT  IN  GRAMS  OF  1  LITER  OF  VARIOUS  GASES  UNDER  STAND- 
ARD CONDITIONS,   AND   BOILING   POINTS  UNDER  PRESSURE   OF 
760  MILLIMETERS 


w 

NAME         OF 
Acetylene  .     .     . 
Air 

EIGHT     BOILING 
LITER     POINT 

1.1621      -83.8° 
2928 

WEIGHT 
NAME         OF  1  LITER 

Hydrogen  chloride  1.6398 
Hydrogen  fluoride  0.893 
Hydrogen  sulfide    1.5392 
Methane     .     .     .0.7168 
Nitric  oxide     .     .  1.3402 
Nitrogen     .     .     .  1.2507 
Nitrous  oxide       .  1.9777 
Oxygen      .     .     .  1.4290 
Sulfur  dioxide     .  2.9266 

BOILING 
POINT 

-82.9° 
+  19.4° 
-61.6° 
-164.0° 
-153.0° 
-195.7° 
-89.8° 
-182.9° 
-8.0° 

Ammonia  .     .     . 
Argon    .... 
Carbon  dioxide   . 
Carbon  monoxide 
Chlorine     .     .     . 
Helium.     .     .     .  ( 
Hydrogen  .     .     . 

).7708      -33.5° 
1.7809    -186.0° 
.9768      -78.2° 
.2504    -190.0° 
U674      -33.6° 
).1782    -268.7° 
X08987-252.70 

DENSITIES  AND  MELTING  POINTS  OF  SOME  COMMON  ELEMENTS 


NAME           I 

)ENSITY 

MELTING 
POINT 

NAME           ] 

DENSITY 

MELTING 
POINT 

Aluminium     .     . 

2.65 

658.7° 

Magnesium    .     . 

1.74 

651.0° 

Antimony      .     . 

6.52 

630.0° 

Manganese     .     . 

8.01 

1260.0° 

Arsenic     .     .     . 

5.73 

850.0° 

Mercury    .     .     . 

13.56 

-38.9° 

Bismuth    .     .     . 

9.80 

271.0° 

Nickel  .... 

8.9 

1452.0° 

Calcium    .     .     . 

1.55 

810.0° 

Phosphorus    .     . 

1.83 

44.0° 

Carbon,  diamond 

3.52 

Platinum  .     .     . 

21.50 

1755.0° 

Carbon,  graphite 

2.30 

>  3600.0° 

Potassium      .     . 

0.862 

62.3° 

Chromium      .     . 

7.3 

1520.0° 

Silicon  .... 

2.3 

1420.0° 

Cobalt       .     .     . 

8.6 

1480.0° 

Silver  .... 

10.5 

960.5° 

Copper      .     .     . 

8.93 

1083.0° 

Sodium      .     .     . 

0.97 

97.5° 

Gold     .... 

19.32 

1063.0° 

Sulfur  (rhombic) 

2.06 

112.8° 

Iron      .... 

7.86 

1530.0° 

Tin  

7.29 

231.9° 

Lead     .... 

11.37 

327.4° 

Zinc      .... 

7.10 

419.4° 

APPENDIX 


561 


SOLUBILITY  OF  VARIOUS  GASES  IN  WATER 


NAME  OF  GAS 


VOLUME  ABSORBED  AT  0°  AND 

UNDER  760  MM.  PRESSURE   BY 

1  LITER  OF  WATER 


Ammonia  .  .  . 
Hydrogen  chloride 
Sulfur  dioxide 
Hydrogen  sulfide  . 
Carbon  dioxide  . 
Oxygen  .... 
Hydrogen  .  .  . 
Nitrogen  .  .  . 


1298.9  liters 
506.0  liters 
79. 79  liters 
4.37  liters 
1.713  liters 
0.0496  liters 
0.0214  liters 
0.0233  liters 


TABLE  OF  SOLUBILITY  OF  VARIOUS  SOLIDS 


SUBSTANCE 

FORMULA 

WEIGHT  DISSOLVED  BY  100  cc.  OF  WATER  AT 

0° 

200 

100° 

Calcium  chloride    . 
Sodium  chloride 
Potassium  nitrate    . 
Copper  sulfate    .     . 
Calcium  sulfate  .     . 
Calcium  hydroxide 

CaCl2 
NaCl 
KN08 
CuS04 
CaS04 
Ca(OH)0 

59.5  g. 
85.70  g. 
13.30  g. 
14.30  g. 
0.759  g. 
0.185  g. 

74.5  g. 
36.0  g. 
31.6  g. 
21.7  g. 
0.203  g. 
0.165  g. 

159.0  g. 
39.80  g. 
246.0  g. 
75.4  g. 
0.162  g. 
0.077  g. 

RELATION  OF  COMMON  UNITS  AND  METRIC  UNITS 

1  pound  (troy)  =  373.24  grams 
1  pound  (avoirdupois)  =  453.59  grams 
1  ounce  (avoirdupois)  =  28.35  grams 
1  United  States  quart  =  0.946  liters 

1  liter  =  1.056  United  States  quarts 
1  meter  =  39.87  inches 
1  centimeter  =  nearly  $  inch 
1  kilogram  =  nearly  2£  pounds  avoirdupois 


INDEX 


Accumulator,  521 

Acetic  acid,  120,  339;  glacial, 
339 

Acetylene,  302 

Acid  anhydrides,  216 

Acid  salts,  preparation  of,  249 

Acids,  179;  binary,  189 ;  character- 
istics of,  179  ;  the  common,  179  ; 
definition  of,  180 ;  dibasic,  249  ; 
ionization  of,  180;  monobasic, 
249 ;  naming  of,  189 ;  organic, 
339;  preparation  of,  225 ;  strength 
of,  188  ;  ternary,  189 ;  tribasic, 
249 ;  undissociated,  181 

Affinity,  15 

Agate,  371 

Agent,  catalytic,  26  ;  dehydrating, 
248  ;  oxidizing,  31  ;  reducing, 
48 

Air,  135 ;  analysis  of,  136 ;  com- 
position of,  135 ;  constancy  of, 
139  ;  constituents  of,  essential  to 
life,  136;  impure,  140;  liquid, 
141  ;  a  mixture,  139  ;  nitrogen  in, 
137 ;  oxygen  in,  136 ;  processes 
tending  to  change  constancy  of, 
137 ;  water  vapor  in,  136 

A.ir  saltpeter,  208 

Alchemists,  20 

Alchemy,  20 

Alcohol,  332  ;  absolute,  334  ;  com- 
mon, 332  ;  denatured,  334  ;  ethyl, 
332;  grain,  332;  methyl,  332; 
wood,  120,  332 

Alcohols,  331 

Alizarin,  336 

Alkali  metals.  396 

Alkalies,  181 

Alkaline  earth  metals,  423 

Allotropic  forms,  112 

Alloys,  364 

Aluminates,  457 


Aluminium,  451 ;  hydrolysis  of 
salts  of,  460 

Aluminium  chloride,  459 

Aluminium  group,  451 

Aluminium  hydroxide,  456 

Aluminium  nitride,  461 

Aluminium  oxide,  456 

Aluminium  silicates,  464 

Aluminium  sulfate,  459 

Alums,  459 

Alundum,  456 

Amalgams,  502 

Amethyst,  371 ;  oriental,  456 

Ammine  salts,  500 

Ammonia,  201 ;  chemical  conduct 
of,  204  ;  composition  of,  205 ;  and 
manufacture  of  ice,  111 ;  prepa- 
ration of,  201  ;  preparation  of, 
from  air,  461  ;  properties  of,  203  ; 
uses  of,  205 

Ammoniacal  liquor,  203,  306 

Ammonium,  205 ;  compounds  of, 
412 

Ammonium  alum,  459 

Ammonium  carbonates,  413 

Ammonium  chloride,  413 

Ammonium  chlorostannate,  515 

Ammonium  hydroxide,  204 

Ammonium  iron  alum,  459,  488 

Ammonium  molybdate,  554 

Ammonium  persulfate,  249 

Ammonium  salts,  205 

Ammonium  sulfate,  413 

Ammonium  sulfides,  413 

Analysis,  73 

Anhydride,  216 ;  chromic,  533 ; 
phosphoric,  352  ;  silicic,  372  ; 
sulfuric,  240  ;  sulfurous,  236 

Anhydrite,  424 

Aniline,  337 

Aniline  dyes,  337 

Anion,  153 


563 


564    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Anode,  151 

Anthracene,  336 

Antimony,  359 ;  acids  of,  360  ;  al- 
loys of,  364  ;  metallic  properties 
of,  361  ;  sulfides  of,  361 

Antimony  chloride,  360 

Antimony  salts,  hydrolysis  of, 
361 

Apatite,  347 

Appendix,  557 

Aqua  ammonia,  204 

Aqua  regia,  212 

Aqueous  vapor  pressure,  57 ;  cor- 
rection for,  58  ;  table  of,  560 

Aragonite,  429 

Arc,  flaming,  554 

Argol,  341 

Argon,  133 

Arrhenius,  Svante  (portrait),  152 

Arsenic,  355 ;  acids  of,  358 ;  Marsh's 
test  for,  357;  oxides  of,  358; 
sulfides  of,  359  ;  white,  358 

Arsenopyrite,  355 

Arsine,  356  ;  combustion  of,  317 

Asbestos,  441, 444, 464  ;  platinized, 
546 

Atmosphere,  135 

Atomic  theory,  88 ;  and  radium, 
540 

Atomic  weights,  91 ;  accurate  de- 
termination of,  287;  from  com- 
bining weights,  286  ;  relation  of 
properties  of  elements  to,  257  ; 
steps  in  determining,  287 

Atoms,  89 ;  replacing  power  of, 
195 

Avogadro,  hypothesis  of,  61,  281 ; 
statue  of,  281 

Babbitt  metal,  364 
Bakelite,  337 
Baking  powders,  460 
Barite,  435 
Barium,  434 
Barium  chloride,  485 
Barium  nitrate,  435 
Barium  oxide,  435 
Barium  peroxide,  82,  435 
Barium  sulf  ate,  435 
Bases,  action  of  halogen  elements 
on,  409  ;  characteristics  of,  181  ; 


the  common,  181 ;  definition  of, 
182  ;  ionization  of,  181 ;  naming 
of,  189 ;  strength  of,  188 

Bauxite,  452 

Benzene,  297,  336 

Benzine,  299 

Benzoic  acid,  337 

Bessemer  converter,  477 

Bessemer  process,  487 

Belts  process,  517 

Birkeland  and  Eyde  process,  208 

Bismuth,  362  ;  alloys  of,  364  ;  com- 
pounds of,  363  ;  hydrolysis  of 
salts  of,  363 

Bismuth  chloride,  363 

Bismuth  oxychloride,  364 

Bismuth  oxynitrate,  364 

Bisque,  465 

Bivalent  elements,  193 

Bleaching,  by  bleaching  powder, 
428;  by  chlorine,  165;  by  hy- 
drogen peroxide,  83  ;  by  ozone, 
114  ;  by  sulfurous  acid,  239 

Bleaching  powder,  427 

Blowpipe,  oxyacetylene,  304 ;  oxy- 
hydrogen,  49 

Blue  printing,  488 

Bluestone,  499 

Boiling  point,  105 ;  of  solutions, 
150 

Bone  ash,  434 

Bone  black,  121 

Borax,  380 

Borax  bead,  381 

Bordeaux  mixture,  499 

Boric  acid,  379 

Bornite,  495 

Boron,  379 

Boyle,  Robert  (portrait),  52 

Boyle's  law,  53 

Brass,  497 

Brazing,  513 

Bread-making,  chemical  changes 
in,  335 

Brick,  vitrified,  465 

Brimstone,  229 

Bromic  acid,  277 

Bromides,  275 

Bromine,  268 

Bronze,  497  ;  aluminium,  454,  497 

Bunsen,  Robert  (portrait),  397 


INDEX 


565 


Bunsen  burner,  314 
Burning,  changes  in   weight  dur- 
ing, 2,  5 ;    questions  suggested 

by,  2 

Butter  fat,  342 
Butyric  acid,  339 
Butyrin,  342 
By-products,  403 

Cadmium,  449  ;  compounds  of,  449 

Caesium,  412 

Calcite,  429 

Calcium,  424 

Calcium  acid  carbonate,  429 

Calcium  acid  sulfite,  434 

Calcium  carbide,  432 

Calcium  carbonate,  428 

Calcium  chloride,  434 

Calcium  cyanide,  433 

Calcium  family,  423 

Calcium  fluoride,  434 

Calcium  hydroxide,  426  ;  action  of 
carbon  dioxide  on,  295 

Calcium  oxide,  425 

Calcium  phosphates,  433,  437 

Calcium  silicates,  434 

Calcium  sulfate,  430 

Calcium  sulfide,  434 

Caliche,  405 

Calomel,  503 

Calorie,  8 

Calorimeter,  8  ;  bomb,  321 

Caramel,  320 

Carbides,  122 

Carbohydrates,  324 

Carbolic  acid,  337 

Carbon,  116 ;  amorphous,  119 ;  crys- 
talline, 116;  cycle  of,  in  nature, 
344  ;  occurrence  of,  116  ;  oxides 
of,  291  ;  properties  of,  121  ;  re- 
tort, 307  ;  uses  of,  122 

Carbon  dioxide,  122  ;  action  of,  on 
calcium  hydroxide,  295  ;  in  air, 
137 

Carbon  disulfide,  249 

Carbon  monoxide,  291 

Carbon  suboxide,  291 

Carbon  tetrachloride,  302 

Carbona,  302 

Carbonates,  294 

Carbonic  acid,  294 ;  salts  of,  294 


Carbonyl  chloride,  292 

Carborundum,  369 

Carnallite,  408 

Carnotite,  535,  554 

Carrel-Dakin  solution,  406 

Casein,  343 

Cassiterite,  511 

Castner  process,  174 

Catalysis,  221  ;  effect  of,  on  speed 

of  reaction,  221 
Catalytic  agent,  26,  221,  546 
Cathode,  151 
Cation,  153 
Caustic  potash,  408 
Caustic  soda,  177 
Caves,  430 
Celestite,  434 
Cell,  Daniell,  500  ;    storage,  521  ; 

electric,  500 
Celluloid,  329 
Cellulose,  328 

Cement,  466 ;  the  setting  of,  466 
Cementite,  471 
Ceresin,  299 
Cerium,  522,  553 
Chalcedony,  371 
Chalcocite,  495 
Chalcopyrite,  495 
Chalk,  French,  445  ;  precipitated, 

429 

Chalybeate  waters,  484 
Chamber  crystals,  244 
Charcoal,  120  ;  animal,  121 
Chemical    action,    14;    conditions 

affecting,  15 
Chemical  affinity,  15 
Chemical  compounds,  15 
Chemical  conduct,  16 
Chemical  energy,  7 
Chemical  reactions,  14 
Chemistry,  domain  of,  21 ;  relation 

of,  to  physics,  1 
Chile  saltpeter,  404 
Chinaware,  465 
Chlorates,  276 
Chlorauric  acid,  551 
Chloric  acid,  276 
Chlorides,  275  ;  definition  of,  162 
Chlorine,  159;  action  as  germicide, 

166;   action  of,   on  compounds 

of  hydrogen,  163;  action  of,  on 


566    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


hydrogen,  163 ;  action  of,  on  met- 
als, 162 ;  action  of,  on  turpentine, 
163 ;  action  of,  on  water,  164  ; 
bleaching  action,  164 ;  family, 
264  ;  occurrence  of,  159  ;  prepa- 
ration of,  159;  properties  of,  162; 
purification  of  water  by,  166 

Chloroform,  302 

Chlorophyll,  441,  470 

Chloroplatinic  acid,  548 

Chromates,  530;  oxidizing  action 
of,  532 

Chrome  alums,  530 

Chrome  yellow,  520,  531 

Chromic  acid,  530 

Chromic  anhydride,  533 

Chromic  chloride,  529 

Chromic  hydroxide,  529 

Chromic  sulfate,  529 

Chromite,  529 

Chromium,  528 

Chromium  salts,  hydrolysis  of, 
530 

Cinnabar,  503 

Citric  acid,  341 

Clay,  464 

Clay  products,  464 

Coal,  119 

Coal  gas,  306 

Coal  oil,  299 

Coal  tar,  306 

Coal-tar  compounds,  336 

Cobalt,  489  ;  compounds  of,  491 

Coin,  gold,  497;  nickel,  497  ;  silver, 
497 

Coke,  119,  307 

Coke  oven,  the  by-product,  307 

Colemanite,  380 

Collodion,  329 

Colloidal  state,  382 

Colloids,  382 ;  coagulation  of,  383  ; 
and  hydrolysis,  385;  industrial 
importance  of,  386 ;  nature  of, 
386 

Combining  weights,  86  ;  law  of,  88 ; 
standards  for,  87 

Combustion,  32 ;  phlogiston  theory 
of,  34  ;  spontaneous,  33 

Composition,  percentage,  93 

Compounds,  definition  of,  16 ;  in- 
soluble, 394 ;  isomeric,  324 


Concentration,  effect  of,  on  speed 

of  reactions,  220 
Concrete,  468  ;  reenf  orced,  468 
Condensation,  heat  of,  106 
Condensed  acids,  373 
Copper,  494;  action  of  nitric  acid 

on,  2i2 ;  action  of  sulfnric  acid 

on,  247  ;  alloys  of,  497 ;  blister, 

496 ;     refining    of,    496 ;    ruby, 

498 

Copperas,  483 
Corn  sirup,  327 
Corrosive  sublimate,  503 
Corundum,  456 
Cotton,  mercerized,  329 
Cotton  fiber,  330 
Coumarin,  337 
Cream  of  tartar,  341 
Cresol,  337 
Critical  point,  106 
Crocoite,  528 
Cryolite,  265,  452 
Crystal  systems,  557 
Crystallization,  methods  of,  148 
Crystallography,  557  ;   systems  of, 

557 
Crystals,    112 ;    growth    of,    559 ; 

structure  of,  558 
Cupellation  of  silver,  504 
Cupric  acetate,  499 
Cupric  ammonia  compounds,  499 
Cupric  bromide,  499 
Cupric  chloride,  499 
Cupric  compounds,  498 
Cupric  nitrate,  499 
Cupric  oxide,  498 
Cupric  sulfate,  498 
Cupric  sulfide,  499 
Cuprite,  495,  498 
Cuprous  chloride,  498 
Cuprous  compounds,  498 
Cuprous  oxide,  498 
Curie,  Madame  (portrait),  637 
Cyanamide,  433 
Cyanides,  296 
Cyanogen,  296 

Dalton,  John  (portrait),  89 
Davy,  Sir  Humphry  (portrait),  173 
Decomposition,  14 
Definite  composition,  law  of,  35,  80 


INDEX 


567 


Deliquescence,  411 

Dewar  flask*,  108 

Dextrin,  327 

Dextrose,  325,  326 

Diamond,  117 

Dichromates,  531  ;  oxidizing  ac- 
tion of,  532 

Dimorphism,  559 

Displacement  series,  191 

Distillation,  68  ;  destructive,  121  ; 
of  water,  68 

Dolomite,  442 

Drununond  light,  50 

Dye,  337 

Dyeing,  458 

Dynamite,  420  ;  gelatin,  421 

Earth's  crust,  composition  of,  19 

Earths,  the  rare,  451,  553 

Efflorescence,  252 

Electrochemical  industries,  391 

Electrochemical  series,  190 

Electrodes,  151 

Electrolysis,  150  ;  of  sodium  chlo- 
ride, 155  ;  of  water,  156 

Electrolyte,  150 

Electrons,  157,  191,  540 

Electroplating,  490 

Electroscope,  536 

Electrotyping,  497 

Elementary  substances,  17 

Elements,  16  ;  classification  of,  254; 
distribution  of,  19  ;  electronega- 
tive, 198  ;  electropositive,  198  ; 
essential  to  life,  19;  metallo- 
acid,  524 ;  molecular  weight  of, 
285  ;  molecules  of,  91 ;  names  of, 
18 ;  number  of,  17  ;  occurrence 
of,  18 ;  physical  state  of,  18 ; 
symbols  of,  17 

Emery,  452,  456 

Emulsions,  387 

Enamels,  377 

Energy,  5  ;  chemical,  7 ;  conserva- 
tion of,  6 ; '  forms  of,  5  ;  and 
matter,  5 ;  measurement  of,  8 ; 
transformations  of,  6 

Epsom  salt,  444 

Equations,  97  ;  molecular,  98 ;  prob- 
lems based  on,  100  ;  and  volumes 
of  gases,  288 


Equilibrium,  222  ;   effect  of  mass 

on,  222  ;  in  solution,  228 
Esters,  341 

Etching  by  hydrofluoric  acid,  268 
Ether,  334 
Eudiometer,  74,  75 
Evaporation,  104 
Explosives,  419 

Families  of  elements,  259 

Fats,  341 

Feldspar,  373,  451 

Fermentation,  333  ;  acetic,  340  ; 
alcoholic,  333  ;  lactic,  326 

Ferric  chloride,  485 

Ferric  hydroxide,  484  ;  colloidal, 
385 

Ferric  nitrate,  488 

Ferric  salts,  484 ;  reduction  of, 
486 

Ferric  sulfate,  488 

Ferrochromium,  528 

Ferromanganese,  481,  525 

Ferromolybdemun,  555 

Ferrosilicon,  368 

Ferrotitanium,  378,  564 

Ferrotungsten,  556 

Ferrous  carbonate,  484 

Ferrous  salts,  483  ;  oxidization  of, 
485 

Ferrous  sulfate,  483 

Ferrous  sulfide,  483 

Ferrovanadium,  554 

Fertilizers,  436,  461 

Filtration,  70 

Fire  damp,  301 

Fire  extinguishers,  126 

Flames,  311;  complex,  315;  con- 
ditions necessary  for,  313  ;  from 
liquids  or  solids,  314  ;  luminosity 
of,  316;  oxidizing,  31 8;  reducing, 
318;  relation  of  gases  to,  311  ; 
simple,  315;  structure  of,  314; 
temperature  of,  317 

Flash  lights,  442 

Fluorapatite,  265 

Fluorides,  267,  275 

Fluorine,  265 

Fluorite,  434 

Fluosilicic  acid,  869 

Flux,  472 


568    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Fogs,  388 

Foods,  343 ;  bleaching  of,  165 ; 
coal-tar  compounds  in,  337 

Formaldehyde,  332 

Formalin,  332 

Formic  acid,  292 

Formulas,  93  ;  facts  expressed  by, 
96 ;  from  percentage  composition, 
94 ;  structural,  197;  from  valence, 
194 

Franklinite,  445 

Freezing  point,  109 ;  of  solutions, 
150 

Fuels,  306  ;  calorific  value  of,  320  ; 
products  of  combustion  of,  313 

Fuller's  earth,  464 

Furnace,  arc,  321 ;  blast,  473  ;  elec- 
tric, 321;  open  hearth,  478; 
puddling,  476;  resistance,  322; 
reverberatory,  495 

Fusion,  heat  of,  110 

Galenite,  618 

Gallium,  261 

Galvanized  iron,  447 ;  rusting  of, 
512 

Gas,  306 ;  coal,  306  ;  natural,  310  ; 
producer,  309  ;  water,  308 

Gas  laws,  the,  52 

Gas  mantles,  312 

Gases,  definition  of,  104  ;  liquefac- 
tion of,  107;  solution  of,  in 
liquids,  144 

Gasoline,  299,  300 

Gay-Lussac  (portrait),  55 ;  law  of 
expansion  of  gases,  55 ;  law  of 
volumes,  280 

Gels,  383 

Gems,  artificial,  456 

German  silver,  489 

Glass,  374  ;  coloring  of,  376  ;  mold- 
ing and  blowing  of,  374  ;  nature 
of,  377  ;  varieties  of,  375 

Glauber's  salt,  401 

Glazing,  465 

Glucose,  327 

Gluten,  327 

Glycerin,  842,  417,  419 

Glyceryl  nitrate,  419,  420 

Gold,  549;  colloidal,  384;  com- 
pounds of,  551 ;  fool's,  484 ; 


mining    of,    549 ;    recovery   of, 

from  copper,  496 ;    refining  of, 

550 

Gold  telluride,  549 
Goldschmidt  process,  391,  454 
Gram-molecular      volume,      284 ; 

weight,  96 
Granite,  373,  464 
Graphite,  118 
Gun  cotton,  329 
Gun  metal,  497 
Gunpowder,  black,  420 ;  smokeless, 

421 
Gypsum,  430 

Haber  process,  203 

Haemoglobin,  470 

Hall,  Charles  M.  (portrait),  452 

Halogens,  264 

Hare,  Robert  (portrait),  48 

Heat,  of  condensation,  106 ;  of 
fusion,  110  ;  measurement  of,  8  ; 
of  neutralization,  184;  of  re- 
action, 99  ;  of  solidification,  110  ; 
of  vaporization,  106 

Helium,  133,  537 

Hematite,  471,  482 

Henry,  law  of,  145 

Human  body,  composition  of,  20 

Humidity,  relative,  105 

Hydrates,  72,  251 

Hydrazine,  206 

Hydrides,  44 

Hydriodic  acid,  274 

Hydrobromic  acid,  272 

Hydrocarbons,  296;  benzene  series 
of,  297  ;  methane  series  of,  297 

Hydrochloric  acid,  40,  168,  170, 
179  ;  salts  of,  275 

Hydrocyanic  acid,  296 

Hydrofluoric  acid,  267  ;  action  of, 
on  silica,  267;  salts  of,  267, 
275 

Hydrogels,  383 

Hydrogen,  38 ;  chemical  conduct 
of,  44  ;  commercial  preparation 
of,  43  ;  preparation  of,  99  ;  prep- 
aration of,  from  acids,  40 ; 
preparation  of,  from  water,  38 ; 
properties  of,  43  ;  standard  for 
valence,  194  ;  uses  of,  50 


INDEX 


569 


Hydrogen  bromide,  270 

Hydrogen  chloride,  167 ;  composi- 
tion of,  169  ;  solubility  of,  170 ; 
volumetric  composition,  170 

Hydrogen  cyanide,  296 

Hydrogen  fluoride,  267 

Hydrogen  iodide,  273 

Hydrogen  nitrate,  209 

Hydrogen  peroxide,  81 ;  prepara- 
tion of,  99 

Hydrogen  sulfate,  246 

Hydrogen  sulfide,  233 ;  chemical 
conduct  of,  234  ;  combustion  of, 
317 

Hydrolysis,  22(5,  301,  303 

Hydronitric  acid,  206 

Hydrosols,  383 

Hydrosulf uric  acid,  234 ;  salts  of, 
235 

Hydroxides,  189 

Hydroxyl  radical,  189 

Hypobromous  acid,  277 

Hypochlorites,  276 

Hypochlorous  acid,  276 

Hypothesis,  Avogadro's,  281 

Ice,  manufacture  of,  110 

Iceland  spar,  429 

Illium,  548 

Ilmenite,  378,  564 

Indicators,  180 

Indigo,  336 

Infusorial  earth,  371 

Inks,  489 

Invar,  490 

lodic  acid,  277 

Iodides,  275 

Iodine,  272  ;  oxides  of,  276  ;  tinc- 
ture of,  273 

lodoform,  273,  302 

lonization,  151  ;  of  acids,  180 ;  of 
bases,  181 ;  extent  of,  154,  187 ; 
and  the  properties  of  solutions, 
154;  of  salts,  186;  theory  of, 
161 

Ions,  152 ;  electrical  charge  on, 
152;  formation  of,  151;  source 
of  charge  on,  167 

Iridium,  549 

Iron,  470  ;  cast,  473  ;  compounds  of, 

s  482 ;  galvanized,  447 ;  oxides  of, 


482  ;  the  rusting  of,  485  ;  varieties 

of,  471  ;  wrought,  476 
Iron  disulfide,  484 
Iron  family,  469 
Iron  ore,  471 
Iron  sulfide,  14 
Isomeric  compounds,  324 
Isomorphism,  659 

Jasper,  371 

Kainite,  437 
Kaolin,  373,  464 
Kerosene,  299 
Kindling  temperature,  33 
Kinetic  theory,  52,  60 
Kipp  apparatus,  42 
Krypton,  133 

Lactic  acid,  326 

Lactose,  325 

Lakes,  458 

Lamp,  blast,  60  ;  tungsten,  556 

Lampblack,  121 

Laughing  gas,  214 

Lavoisier  (portrait),  frontispiece 

Law,  periodic,  267  ;  of  Boyle,  52  ; 
of  Charles,  55 ;  of  combining  vol- 
umes, 280;  of  combining  weights, 
88  ;  of  conservation  of  energy,  6  ; 
of  conservation  of  matter,  9  ;  of 
definite  composition,  35,  80  ;  of 
Gay-Lnssac,  55,  280;  of  Henry, 
145  ;  of  multiple  proportion,  83 

Laws,  the  gas,  52 

Laws,  meaning  of,  in  science,  60 

Laws  of  Raoult,  285 

Lead,  575  ;  desilverization  of,  516  ; 
hard,  516 ;  oxides  of,  617 ;  a 
product  of  radioactivity,  638 ; 
red,  518  ;  softening  of,  616; 
sugar  of,  339,  520  ;  white,  619 

Lead  acetate,  520 

Lead  arsenate,  359 

Lead  carbonate,  518 

Lead  chloride,  520 

Lead  chromate,  520 

Lead  nitrate,  520 

Lead  peroxide,  518 

Lead  sulfate,  520 

Lead  sulfide,  518 


570    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Leblano,  Nicolas  (statue),  403 

Leblanc  process,  401 

Levulose,  325 

Liebig,  Justus  (portrait),  436 

Liebjg  condenser,  68 

Life,  elements  essential  to,  19 

Liine,  425  ;  air-slaked,  426  ;  slaked, 
426  ;  slaking  of,  425  ;  superphos- 
phate of,  437 

Limekiln,  426 

Limelight,  50,  425 

Lime-nitrogen,  433 

Limestone,  428  ;  doloinitic,  443 

Lime-sulfur  spray,  232 

Limewater,  426 

Liming  soils,  437 

Limonite,  471 

Liquefaction  of  gases,  107 

Liquid-air  machine,  108 

Liquids,  definition  of,  104  ;  under- 
cooled,  109  ;  vapor  pressure  of, 
105 

Litharge,  518 

Lithium,  412 

Lithopone,  448 

Loadstone,  483 

Lunar  caustic,  506 

Lye,  177 

Magnalium,454 
Magnesia,  442 
Magnesia  alba,  443 
Magnesia  usta,  442 
Magnesite,  442 
Magnesium,  441 
Magnesium  carbonate,  442 
Magnesium  chloride,  444 
Magnesium  family,  440 
Magnesium  hydroxide,  442 
Magnesium  nitride,  204 
Magnesium  oxide,  442 
Magnesium  silicate,  444 
Magnesium  sulfate,  444 
Magnetite,  471,  483 ;  structure  of, 

197 

Malachite,  495 
Malt,  334 
Manganates,  526 
Manganese,  524  ;  oxides  of,  525 
Manganese  tetrachloride,  169 
Manganic  acid,  526 


Manganous  carbonate,  525 

Manganous  chloride,  525 

Manganous  hydroxide,  525 

Manganous  sulfate,  525 

Manganous  sulfide,  525 

Marble,  429 

Marsh  gas,  301 

Massicot,  518 

Matches,  350 

Matte,  496 

Matter,  5  ;  and  energy,  5 ;  law 
of  conservation  of,  9  ;  and  its 
properties,  9 ;  states  of,  104  ; 
varieties  of,  12 

Meerschaum,  444,  464 

Melaconite,  495 

Melting  point,  110 

Mendel6eff  (portrait),  255 

Meniscus,  107 

Mercuric  chloride,  503 

Mercuric  fulminate,  504 

Mercuric  nitrate,  504 

Mercuric  oxide,  502 ;  decomposi- 
tion of,  25,  221 

Mercuric  sulfate,  503 

Mercuric  sulfide,  503 

Mercurous  chloride,  503 

Mercurous  nitrate,  504 

Mercurous  sulfate,  504 

Mercury,  501  ;  freezing  of,  by  car- 
bon dioxide,  125 

Metaboric  acid,  380 

Metallo-acid  elements,  524 

Metallurgy,  390 

Metals,  172,  254 ;  definition  of,  389  ; 
extraction  of,  390  ;  occurrence  of, 
in  nature,  390  ;  preparation  of 
compounds  of,  392 ;  properties 
of,  389 

Metaphosphoric  acid,  354 

Metasilicic  acid,  372 

Metastannic  acid,  513 

Meteorites,  469 

Methane,  301  ;  combustion  of,  317 

Mica,  373,  464 

Microcosmic  salt,  353 

Milk,  387 

Minerals,  390 

Miner's  safety  lamp,  301 

Mining,  hydraulic,  550 ;  placer, 
549 ;  vein,  550 


INDEX 


5T1 


Minium,  618 

Mixtures,  20 

Moissan  (portrait),  266 

Molasses,  325 

Molecular  weights,  279 ;  determina- 

.    tion  of,  284  ;  from  weight  of  one 

liter,  283 
Molecules,    60 ;    and    atoms,    90 ; 

of   elements,  91 ;    structure  of, 

196 

Molybdenum,  554 
Monazite  sand,  312 
Monel  metal,  490 
Mordant,  458 
Mortar,  427 
Moth  balls,  336 

Multiple  proportion,  law  of,  83 
Muriatic  acid,  168 

Naphtha,  299 

Naphthalene,  336 

Nascent  state,  167 

Natural  gas,  310 

Nature,  changes  in,  2 

Neon,  133 

Neutralization,  182  ;  a  definite  act, 
183  ;  heat  of,  184 

Nickel,  489 ;  compounds  of,  491 ; 
recovery  of,  from  copper, 
496 

Nickel  coinage,  489 

Nickel  plating,  490 

Nickel  steel,  489 

Niton,  537 

Nitrates,  213 

Nitric  acid,  179,  206  ;  action  of,  on 
metals,  210;  chemical  conduct  of, 
209  ;  decomposition  of,  by  heat, 
210;  oxidizing  action  of,  210; 
preparation  of,  207 ;  properties 
of,  209 ;  salts  of,  213 

Nitric  oxide,  215 

Nitrites,  213 

Nitrobenzene,  337 

Nitrocellulose,  329,  420 

Nitrogen,  128 ;  assimilation  of,  by 
plants,  132  ;  compounds  of,  200  : 
determination  of,  in  air,  137 : 
oxides  of,  214 ;  preparation  of, 
from  air,  129 ;  preparation  of, 
from  compounds,  130 ;  properties 


of,  131 ;  the  utilization  of  atmos- 
pheric, 461 

Nitrogen  dioxide,"  216 
Nitrogen  pentoxide,  214,  216 
Nitrogen  tetroxide,  216 
Nitrogen  trioxide,  214,  216 
Nitroglycerin,  420 
Nitrosyl  sulfuric  acid,  244 
Nitrous  acid,  213 
Nitrous  oxide,  214 
Nonmetals.  172,  264 

Oil,  corn,  327 ;  of  vitriol,  246 

Oils,  341 ;  changing  of,  into  fats, 
343  ;  cracking  of,  300 

Oleic  acid,  341,  342 

Olein,  342 

Oleomargarine,  342 

Onyx,  371 

Opal,  371 

Open-hearth  process,  478 

Ores,  390 

Organic  acids,  339 

Organic  chemistry,  291 

Orpiment,  355 

Orthophosphates,  353 

Orthosilicic  acid,  372 

Osmic  acid,  549 

Osmium,  544 

Osmium  tetroxide,  549 

Oxalic  acid,  292 

Oxidation,  31;  and  combustion, 
32 ;  definition,  486 ;  products 
of,  31 ;  relation  of,  to  reduction, 
48 

Oxygen,  24  ;  chemical  conduct  of, 
29 ;  determination  of,  in  air, 
136  ;  historical,  24  ;  importance 
of,  33  ;  occurrence  of,  24 ;  prepa- 
ration of,  99 ;  preparation  of, 
from  air,  28 ;  preparation  of,  in 
laboratory,  27;  preparation  of, 
from  mercuric  oxide,  25 ;  prepa- 
ration of,  from  potassium  chlo- 
rate, 25 ;  preparation  of,  from 
sodium  peroxide,  27 ;  preparation 
of,  from  water,  25;  properties 
of,  29 ;  as  standard  for  molecular 
weights,  282 

Oxy  salts,  362 

Ozone,  113 


572    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Paints,  520 

Palau,  548 

Palladium,  548 

Palmitic  acid,  342 

Palmitin,  342 

Paper,  330 

Paraffin,  299 

Paris  green,  359 

Parkes  process,  516 

Perchlorates,  277 

Perchloric  acid,  277 

Periodic  acid,  277 

Periodic  classification,  255 

Periodic  families,  259 

Periodic  group,  259 

Periodic  law,  257  ;  value  of,  260 

Periodic  table,  258 

Permanganates,     526 ;      oxidizing 

properties  of,  527 
Permanganic  acid,  526 
Persulfates,  249 
Persulfuric  acid,  249 
Petroleum,  298 
Pewter,  512 
Phenol,  337 

Philosophers1  stone,  20,  347 
Phlogiston,  34 
Phosgene,  292 
Phosphates,  353 
Phosphine,  351 

Phosphonium  compounds,  351 
Phosphorescence,  349 
Phosphoric     acid,     353 ;     glacial, 

354 

Phosphorite,  347,  434 
Phosphorous  acid,  352 
Phosphorus,  347;  acids  of,  352; 

chlorides  of,  355 ;   hydrides  of, 

351  ;   oxides  of,  352  ;   red,  349  ; 

sulfides  of,  355  ;  white,  348 
Phosphorus  family,  346 
Phosphorus  pentoxide,  352 
Phosphorus  trioxide,  352 
Photography,  508 
Picric  acid,  421 
Pink  salt,  515 
Pitchblende,  535 
Plaster,  427 
Plaster  of  Paris,  430 
Platinum,  545 ;  application  of,  547; 

as  a  catalytic  agent,  241,  546 ; 


colloidal,  384 ;  compounds  of, 
548 ;  substitutes  for,  547 ;  use 
of,  as  catalyzer,  243 

Platinum  metals,  544 

Point,  boiling,  105  ;  critical,  106 ; 
freezing,  109;  melting,  110 

Porcelain,  465 

Portland  cement,  466 

Potash,  caustic,  408 

Potassium,  406 ;  and  plant  life, 
407 ;  sources  of,  408 

Potassium  alum,  459 

Potassium  bromide,  409 

Potassium  carbonate,  411 

Potassium  chlorate,  410 

Potassium  chloride,  409 

Potassium  chromate,  531 

Potassium  chrome  alum,  459,  529, 
533 

Potassium  cyanide,  411 

Potassium  dichromate,  531 

Potassium  f  erricyanide,  488 

Potassium  ferrocyanide,  487 

Potassium  hydroxide,  408 

Potassium  iodide,  410 

Potassium  manganate,  526 

Potassium  nitrate,  410 

Potassium  permanganate,  526 ; 
use  of,  in  preparing  chlorine, 
161 

Potassium  silicate,  373 

Potassium  sulfate,  411 

Potassium  sulfite,  411 

Pottery,  white,  465 

Precipitates,  225 

Preservatives,  335 

Pressure,  critical,  107;  of  atmos- 
phere, 53  ;  standard,  53 

Priestley,  Joseph  (portrait),  24 

Producer  gas,  309 

Properties,  definition  of,  9 

Proteins,  343 

Prussian  blue,  487 

Prussiate  of  potash,  red,  488  ;  yel- 

•  low,  487 

Prussic  acid,  296 

Pyrite,  484 

Pyrolusite,  525 

Pyrophosphoric  acid,  354 

Pyrosulfuric  acid,  249 

Pyrrhotite,  483 


INDEX 


573 


Quadrivalent  elements,  193 
Quartz,  370 
Quicklime,  425 

Radical,  glyceryl,  342  ;  hydroxyl, 
189 

Radicals,  188 ;  replacing  power  of, 
195  ;  valence  of,  195 

Radioactivity,  535 

Radiograph,  53(5 

Radium,  1 7,  536  ;  and  the  atomic 
conception,  540 ;  discovery  of, 
536  ;  and  disease,  541 ;  disintegra- 
tion of,  537;  energy  of,  539;  origin 
of,  539 ;  quantity  of  available, 
537 ;  radiation  from,  538 

Ramsay,  Sir  William  (portrait),  133 

Raoult,  laws  of,  285 

Reactions,  completion  of,  in  solu- 
tion, 224 ;  definition  of,  14 ;  endo- 
thermic,  50 ;  exothermic,  50 ; 
factors  affecting  speed  of,  219; 
flame,  414  ;  heat  of,  99  ;  reversi- 
ble, 221 ;  speed  of,  219 

Realgar,  355 

Reduction,  47,  48,  48(5 ;  relation  of, 
to  oxidation,  48 

Rhodium,  544 

Rhotanium,  548 

Rifle  bullets,  490 

Roasting,  390 

Rouge,  482 

Rubidium,  412 

Ruby,  456 

Ruthenium,  544 

Rutile,  378,  554 


Saccharine,  337 

Sal  ammoniac,  413 

Sal  soda,  403 

Salt,  common,  399  ;  rock,  399 

Saltpeter,  410 

Salts,  184 ;   acid,  186 ;   basic,  187, 

362  ;   definition  of,  185  ;  ioniza- 

tion   of,  186  ;    naming  of,  190  ; 

normal,  186  ;  oxy,  362 ;  relation 

of,  to  acids,  185 
Sand,  371 
Sandstone,  371 
Saponificat.ion,  417 
Sapphire,  466 


Saturation,  104,  148 

Scale,  boiler,  443 

Scheele  (portrait),  128 

Sciences,  the  natural,  1 

Selenium,  251,  555 

Serpentine,  441,  444,  464 

Sewage-disposal  plant,  34 

Shot,  356 

Siderite,  471,  484 

Silica,  370 

Silicates,  373 

Silicic  acids,  372;  condensed,  373; 
simple,  372 

Silicides,  369 

Silicon,  367 

Silicon  dioxide,  370 ;  action  of 
hydroflupric  acid  on,  372 

Silicon  fluoride,  369 

Silicon  hydride,  368 

Silk,  artificial,  329 

Silk  fiber,  330 

Silver,  504  ;  coin,  497  ;  compounds 
of,  505  ;  German,  497  ;  recovery 
of,  from  copper,  496 ;  sterling, 
506 

Silver  bromide,  507 

Silver  chloride,  507 

Silver  chromates,  507 

Silver  cyanides,  507 

Silver  iodide,  507 

Silver  nitrate,  506 

Silver  nitrite,  507 

Silver  plating,  506 

Silver  sulfate,  507 

Silver  sulfide,  507 

Slag,  472 

Smalt,  491 

Smithsonite,  445 

Smoke  prevention,  319 

Soap,  416 ;  cleansing  action  of, 
419 ;  composition  of,  416 ;  manu- 
facture of,  417  ;  properties  of, 
418  ;  salting  out  of,  417  ;  scour- 
ing, 371  ;  varieties  of,  418 

Soapstone,  444 

Soda,  403 ;  baking,  404 ;  bicarbon- 
ate of,  404 ;  caustic,  175 ;  sal, 
403  ;  washing,  403 

Soda  ash,  403 

Sudamide,  405 

Sodium,  172;  compounds  of,  398 


574    AX  ELEMENTARY  STUDY  OF  CHEMISTRY 


Sodium  benzoate,  336,  887 

Sodium  bromide,  400 

Sodiu'm  oarbonate,  401 

Sodium  chlorate,  406 

Sodium  chloride,  399 ;  ionization 
of,  155 

Sodium  chromate,  531 

Sodium  cyanide,  405,  433 

Sodium  dichromate,  531 

Sodium  diuranate,  535 

Sodium  ferrocyanide,  487 

Sodium  hydroxide,  175 

Sodium  hydrogen  carbonate,  404 

Sodium  hypochlorite,  406 

Sodium  hyposulfite,  401 

Sodium  iodide,  400 

Sodium  nitrate,  404 

Sodium  permanganate,  527 

Sodium  peroxide,  27,  398 

Sodium  phosphates,  353,  405 

Sodium  silicate,  373 

Sodium  sulfate,-400 

Sodium  sulfite,  401 

Sodium  thiosulfate,  401 

Solder,  512 

Soldering,  513 

Solids,  109 ;  amorphous,  109 ;  crys- 
talline, 109  ;  definition  of,  104  ; 
evaporation  of,  105 ;  solutions 
of,  146 

Sols,  393  ;  preparation  of,  384 

Solubility,  of  gases,  144  ;  of  solids, 
147 

Solubility  curves,  149 

Solute,  143 

Solutions,  143 ;  boiling  point  of, 
150  ;  definition  of,  143  ;  equilib- 
rium in,  223;  freezing  point  of, 
160 ;  of  gases  in  liquids,  144 ; 
gram-molecular,  147 ;  molar,  147 ; 
properties  of,  150 ;  saturated, 
147;  of  solids  in  liquids,  146; 
super-saturated,  149 

Solvay  process,  402 

Solvent,  143 

Spelter,  445 

Sphalerite,  445 

Spinel,  457 

Spirits  of  hartshorn,  201 

Spray,  lime-sulfur,  232 

Standard  conditions,  56 


Stannatee,  f>13 

Stannic  acid,  513 

Stannic  chloride,  514 

Stannic  oxide,  513 

Stannic  sultide,  515 

Stannous  chloride,  514 

Stannous  sulfide,  515 

Starch,  327 

Stassf  urt  salts,  407 

State,  nascent,  167  ;  native,  18 

Stearic  acid,  342 

Stearin,  342 

Steel,  476  ;  alloy  of,  481  ;  crucible, 
479  ;  electrothermal  metallurgy 
of,  480 ;  hardening  of,  480 ; 
properties  of,  480 ;  tempering 
of,  480 ;  tool,  479  ;  vanadium, 
654 

Steel  purifiers,  482 

Steel  scavengers,  482 

Stellite,  548 

Stibine,  360 

Stibnite,  359 

Strontianite,  434 

Strontium  hydroxide,  435 

Strontium  nitrate,  436 

Structural  formulas,  197 

Sucrose,  325 

Sugar,  325;  beet,  325;  cane,  325; 
grape,  326  ;  invert,  325 ;  milk,  326 

Sugar  of  lead,  339,  520 

Sulfates,  248 

Sulfides,  235  ;  insoluble,  235  ;  solu- 
ble, 235 

Sulfites,  240 

Sulfur,  amorphous,  231  ;  chemical 
conduct  of,  232  ;  extraction  of, 
229;  flowers  of,  229;  mono- 
clinic,  230  ;  occurrence  of,  228  ; 
oxides  of,  236 ;  plastic,  231 ; 
rhombic,  230;  roll,  229;  varie- 
ties of,  230 

Sulfur  dioxide,  236 

Sulfur  trioxide,  240 

Sulfuric  acid,  40,  179,  242  ;  acid 
properties  of,  247  ;  action  of,  on 
metals,  247  ;  action  of,  on  salts, 
248  ;  action  of,  on  water,  248  ; 
chamber  process  for,  243  ;  con- 
tact process  for,  242  ;  manufac- 
ture of,  242  ;  oxidizing  properties 


INDEX 


575 


of,  247  ;  properties  of,  246  ;  salts 
of,  248  ;  structure  of,  196 

Sulf urous  acid,  239  ;  salts  of,  240 

Superphosphate  of  lime,  437 

Supersaturation,  149 

Symbol,  17 

Synthesis,  73 

Talc,  441,  444,  464 

Tannic  acid,  480 

Tartaric  acid,  341 

Tellurium,  251,  553 

Temperature,  absolute  scale  of, 
54  ;  centigrade  scale  of,  54  ;  criti- 
cal. 107  ;  effect  of,  on  speed  of 
reactions,  219;  kindling,  33; 
variation  of  volume  with,  54 

Tetraboric  acid,  380 

Textile  fibers,  330 

Thallium,  451 

Theory,  60  ;  atomic,  88  ;  forming 
a,  60 ;  of  ionization,  151  ;  the 
kinetic,  52,  60;  value  of  a,  61 

Thermite,  455 

Thermite  welding,  455 

Thermoelectric  processes,  391 

Thermos  bottles,  108 

Thorium,  522 ;  disintegration  series 
for,  541  ;  radioactivity  of,  541 

Thorium  oxide,  554 

Tin,  511  ;  block,  511  ;  chlorides  of, 
514;  crystals,  514;  oxymuriate 
of,  514 

Tin  plate,  512  ;  rusting  of,  512 

Titanium,  378,  554  ;  compounds  of, 
379 

Titanium  nitride,  378 

Toluene,  297,  336 

Topaz,  456 

Townsend  cell,  176 

Transition  point,  231 

Trinitrotoluene,  421 

Trivalent  elements,  193 

Tungsten,  555 

Type  metal,  364 

Units  and  abbreviations,  10 
Univalent  elements,  193 
Uraninite,  536 

Uranium,  535;  disintegration  scries 
for,  541  ;  radioactivity  of,  535 


Uranium  steel,  535 
Uranyl.  nitrate,  535 

Valence,  193  ;  cause  of,  197  ;  defi- 
nition of,  193 ;  formulas  from,  194 ; 
standards  for,  194 ;  and  structure 
of  molecules,  196  ;  table  of,  198 ; 
variable,  196 ;  variety  in,  193 

Vanadium,  554  • 

Vanillin,  337 

Vapor,  saturated,  104 

Vapor  pressure,  57  ;  of  hydrates, 
252  ;  of  liquids,  105 ;  measure- 
ment of,  59 

Vaporization,  heat  of,  106 

Vaseline,  299 

Venetian  red,  482 

Ventilation,  140 

Vermilion,  503 

Vinegar,  340 

Vitriol,  blue,  251,  499  ;  green,  483  ; 
oil  of,  246  ;  white,  448 

Volumes,  law  of,  280 

Water,  64  ;  analysis  of,  73  ;  chem- 
ical conduct  of,  72  ;  city  filtra- 
tion of,  70  ;  composition  of,  75  ; 
composition  of,  natural,  65 ; 
composition  of,  by  volume,  78  ; 
detection  of  impurities  in,  66  ; 
distillation  of,  68-;  distilled,  68  ; 
Dumas's  synthesis  of,  77  ;  elec- 
trical decomposition  of,  16  ;  elec- 
trolysis of,  156  ;  exact  composi- 
tion of,  73 ;  hard,  65,  431 ;  heat  of 
formation  of,  100 ;  and  health,  66 ; 
microorganisms  in,  66  ;  mineral, 
65 ;  mineral  analysis  of,  67 ; 
Morley's  synthesis  of,  77 ;  oc- 
currence of,  64 ;  permanent 
hardness  of,  431  ;  properties  of, 
72 ;  purification  of,  68,  457 ; 
purification  of,  by  boiling,  70 ; 
purification  of,  by  chlorine,  166  ; 
purification  of,  by  filtration,  70  ; 
purification  of,  by  ozone,  114 ; 
sanitary  analysis  of,  67 ;  self- 
purification  of,  71  ;  soft,  65  ; 
softening  of,  431  ;  synthesis  of, 
74  ;  temporary  hardness  of,  431 ; 
and  typhoid  fever,  66 


576    AN  ELEMENTARY  STUDY  OF  CHEMISTRY 


Water  gas,  308 

Water  glass,  373 

Water  vapor,  determination  of,  in 
air,  137 

Weight  of  a  liter  of  gas,  288 

Weights,  atomic,  91  ;  combining, 
86  ;  formula,  96  ;  gram-atomic, 
97  ;  gram-molecular,  96  ;  sym- 
bol, 97 

Willemite,  445 

Witherite,  434 

Wollastonite,  373,  424 

Wood,  preservation  of,  448 

Wood's  metal,  365 

Wool  fiber,  330 


Xenon,  133 
Yeast,  333 

Zinc,  445  ;  action  of  nitric  acid  on, 
211 ;  granulated,  446 ;  mossy,  446 
Zinc  chloride,  448 
Zinc  oxide,  447 
Zinc  sulfate,  448 
Zinc  sulfide,  448 
Zinc  white,  447 
Zincite,  445 
Zircite,  554 
Zirconalba,  554 
Zirconium  oxide,  554 


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